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Mark Rosengartens Amazing Chemistry Powerpoint Presentation

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6) Kinetics and Thermodynamics (PE Diagrams, etc. ... Bohr Model. Quantum-Mechanical Model (c) 2006, Mark Rosengarten. Thompson Model ... – PowerPoint PPT presentation

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Title: Mark Rosengartens Amazing Chemistry Powerpoint Presentation


1
Mark Rosengartens Amazing Chemistry Powerpoint
Presentation!
  • Aligned to the New York State Standards and Core
    Curriculum for The Physical Setting-Chemistry
  • Can be used in any high-school chemistry class!
  • Please give the link to this file to your
    chemistry students! www.markrosengarten.com
  • Enjoy it!!! A LOT of work has gone into
    bringing you this work, so please credit me when
    you use it!

2
Outline for Review
  • 1) The Atom (Nuclear, Electron Config)
  • 2) Matter (Phases, Types, Changes)
  • 3) Bonding (Periodic Table, Ionic, Covalent)
  • 4) Compounds (Formulas, Reactions, IMAFs)
  • 5) Math of Chemistry (Formula Mass, Gas Laws,
    Neutralization, etc.)
  • 6) Kinetics and Thermodynamics (PE Diagrams,
    etc.)
  • 7) Acids and Bases (pH, formulas, indicators,
    etc.)
  • 8) Oxidation and Reduction (Half Reactions,
    Cells, etc.)
  • 9) Organic Chemistry (Hydrocarbons, Families,
    Reactions)

3
The Atom
  • 1) Nucleons
  • 2) Isotopes
  • 3) Natural Radioactivity
  • 4) Half-Life
  • 5) Nuclear Power
  • 6) Electron Configuation
  • 7) Development of the Atomic Model

4
Nucleons
  • Protons 1 each, determines identity of
    element, mass of 1 amu, determined using atomic
    number, nuclear charge
  • Neutrons no charge, determines identity of
    isotope of an element, 1 amu, determined using
    mass number - atomic number (amu atomic mass
    unit)
  • 3216S and 3316S are both isotopes of S
  • S-32 has 16 protons and 16 neutrons
  • S-33 has 16 protons and 17 neutrons
  • All atoms of S have a nuclear charge of 16 due
    to the 16 protons.

5
Isotopes
  • Atoms of the same element MUST contain the same
    number of protons.
  • Atoms of the same element can vary in their
    numbers of neutrons, therefore many different
    atomic masses can exist for any one element.
    These are called isotopes.
  • The atomic mass on the Periodic Table is the
    weight-average atomic mass, taking into account
    the different isotope masses and their relative
    abundance.
  • Rounding off the atomic mass on the Periodic
    Table will tell you what the most common isotope
    of that element is.

6
Weight-Average Atomic Mass
  • WAM (( A of A/100) X Mass of A) (( A of
    B/100) X Mass of B) …
  • What is the WAM of an element if its isotope
    masses and abundances are
  • X-200 Mass 200.0 amu, abundance 20.0
  • X-204 Mass 204.0 amu, abundance 80.0
  • amu atomic mass unit (1.66 10-27
    kilograms/amu)

7
Most Common Isotope
  • The weight-average atomic mass of Zinc is 65.39
    amu. What is the most common isotope of Zinc?
    Zn-65!
  • What are the most common isotopes of
  • Co Ag
  • S Pb
  • FACT one atomic mass unit (1.66 10-27
    kilograms) is defined as 1/12 of the mass of an
    atom of C-12.
  • This method doesnt always work, but it usually
    does. Use it for the Regents exam.

8
Natural Radioactivity
  • Alpha Decay
  • Beta Decay
  • Positron Decay
  • Gamma Decay
  • Charges of Decay Particles
  • Natural decay starts with a parent nuclide that
    ejects a decay particle to form a daughter
    nuclide which is more stable than the parent
    nuclide was.

9
Alpha Decay
  • The nucleus ejects two protons and two neutrons.
    The atomic mass decreases by 4, the atomic number
    decreases by 2.
  • 23892U ?

10
Beta Decay
  • A neutron decays into a proton and an electron.
    The electron is ejected from the nucleus as a
    beta particle. The atomic mass remains the same,
    but the atomic number increases by 1.
  • 146C ?

11
Positron Decay
  • A proton is converted into a neutron and a
    positron. The positron is ejected by the
    nucleus. The mass remains the same, but the
    atomic number decreases by 1.
  • 5326Fe ?

12
Gamma Decay
  • The nucleus has energy levels just like
    electrons, but the involve a lot more energy.
    When the nucleus becomes more stable, a gamma ray
    may be released. This is a photon of high-energy
    light, and has no mass or charge. The atomic
    mass and number do not change with gamma. Gamma
    may occur by itself, or in conjunction with any
    other decay type.

13
Charges of Decay Particles
14
Half-Life
  • Half life is the time it takes for half of the
    nuclei in a radioactive sample to undergo decay.
  • Problem Types
  • Going forwards in time
  • Going backwards in time
  • Radioactive Dating

15
Going Forwards in Time
  • How many grams of a 10.0 gram sample of I-131
    (half-life of 8 days) will remain in 24 days?
  • HL t/T 24/8 3
  • Cut 10.0g in half 3 times 5.00, 2.50, 1.25g

16
Going Backwards in Time
  • How many grams of a 10.0 gram sample of I-131
    (half-life of 8 days) would there have been 24
    days ago?
  • HL t/T 24/8 3
  • Double 10.0g 3 times 20.0, 40.0, 80.0 g

17
Radioactive Dating
  • A sample of an ancient scroll contains 50 of the
    original steady-state concentration of C-14. How
    old is the scroll?
  • 50 1 HL
  • 1 HL X 5730 y/HL 5730y

18
Nuclear Power
  • Artificial Transmutation
  • Particle Accelerators
  • Nuclear Fission
  • Nuclear Fusion

19
Artificial Transmutation
  • 4020Ca _____ ----- 4019K 11H
  • 9642Mo 21H ----- 10n _____
  • Nuclide Bullet -- New Element Fragment(s)
  • The masses and atomic numbers must add up to be
    the same on both sides of the arrow.

20
Particle Accelerators
  • Devices that use electromagnetic fields to
    accelerate particle bullets towards target
    nuclei to make artificial transmutation possible!
  • Most of the elements from 93 on up (the
    transuranium elements) were created using
    particle accelerators.
  • Particles with no charge cannot be accelerated by
    the charged fields.

21
Nuclear Fission
  • 23592U 10n ? 9236Kr 14156Ba 3 10n
    energy
  • The three neutrons given off can be reabsorbed by
    other U-235 nuclei to continue fission as a chain
    reaction
  • A tiny bit of mass is lost (mass defect) and
    converted into a huge amount of energy.

22
Chain Reaction
23
Nuclear Fusion
  • 21H 21H ? 42He energy
  • Two small, positively-charged nuclei smash
    together at high temperatures and pressures to
    form one larger nucleus.
  • A small bit of mass is destroyed and converted
    into a huge amount of energy, more than even
    fission.

24
Electron Configuration
  • Basic Configuration
  • Valence Electrons
  • Electron-Dot (Lewis Dot) Diagrams
  • Excited vs. Ground State
  • What is Light?

25
Basic Configuration
  • The number of electrons is determined from the
    atomic number.
  • Look up the basic configuration below the atomic
    number on the periodic table. (PEL principal
    energy level shell)
  • He 2 (2 e- in the 1st PEL)
  • Na 2-8-1 (2 e- in the 1st PEL, 8 in the 2nd and
    1 in the 3rd)
  • Br 2-8-18-7 (2 e- in the 1st PEL, 8 in the 2nd,
    18 in the 3rd and 7 in the 4th)

26
Valence Electrons
  • The valence electrons are responsible for all
    chemical bonding.
  • The valence electrons are the electrons in the
    outermost PEL (shell).
  • He 2 (2 valence electrons)
  • Na 2-8-1 (1 valence electron)
  • Br 2-8-18-7 (7 valence electrons)
  • The maximum number of valence electrons an atom
    can have is EIGHT, called a STABLE OCTET.

27
Electron-Dot Diagrams
  • The number of dots equals the number of valence
    electrons.
  • The number of unpaired valence electrons in a
    nonmetal tells you how many covalent bonds that
    atom can form with other nonmetals or how many
    electrons it wants to gain from metals to form an
    ion.
  • The number of valence electrons in a metal tells
    you how many electrons the metal will lose to
    nonmetals to form an ion. Caution May not work
    with transition metals.
  • EXAMPLE DOT DIAGRAMS

28
Example Dot Diagrams
Carbon can also have this dot diagram, which
it has when it forms organic compounds.
29
Excited vs. Ground State
  • Configurations on the Periodic Table are ground
    state configurations.
  • If electrons are given energy, they rise to
    higher energy levels (excited state).
  • If the total number of electrons matches in the
    configuration, but the configuration doesnt
    match, the atom is in the excited state.
  • Na (ground, on table) 2-8-1
  • Example of excited states 2-7-2, 2-8-0-1, 2-6-3

30
What Is Light?
  • Light is formed when electrons drop from the
    excited state to the ground state.
  • The lines on a bright-line spectrum come from
    specific energy level drops and are unique to
    each element.
  • EXAMPLE SPECTRUM

31
EXAMPLE SPECTRUM
This is the bright-line spectrum of hydrogen.
The top numbers represent the PEL (shell) change
that produces the light with that color and the
bottom number is the wavelength of the light (in
nanometers, or 10-9 m). No other element has
the same bright-line spectrum as hydrogen, so
these spectra can be used to identify elements or
mixtures of elements.
32
Development of the Atomic Model
  • Thompson Model
  • Rutherford Gold Foil Experiment and Model
  • Bohr Model
  • Quantum-Mechanical Model

33
Thompson Model
  • The atom is a positively charged diffuse mass
    with negatively charged electrons stuck in it.

34
Rutherford Model
  • The atom is made of a small, dense, positively
    charged nucleus with electrons at a distance, the
    vast majority of the volume of the atom is empty
    space.

Alpha particles shot at a thin sheet of
gold foil most go through (empty space).
Some deflect or bounce off (small
charged nucleus).
35
Bohr Model
  • Electrons orbit around the nucleus in energy
    levels (shells). Atomic bright-line spectra was
    the clue.

36
Quantum-Mechanical Model
  • Electron energy levels are wave functions.
  • Electrons are found in orbitals, regions of space
    where an electron is most likely to be found.
  • You cant know both where the electron is and
    where it is going at the same time.
  • Electrons buzz around the nucleus like gnats
    buzzing around your head.

37
Matter
  • 1) Properties of Phases
  • 2) Types of Matter
  • 3) Phase Changes

38
Properties of Phases
  • Solids Crystal lattice (regular geometric
    pattern), vibration motion only
  • Liquids particles flow past each other but are
    still attracted to each other.
  • Gases particles are small and far apart, they
    travel in a straight line until they hit
    something, they bounce off without losing any
    energy, they are so far apart from each other
    that they have effectively no attractive forces
    and their speed is directly proportional to the
    Kelvin temperature (Kinetic-Molecular Theory,
    Ideal Gas Theory)

39
Solids
The positive and negative ions alternate in the
ionic crystal lattice of NaCl.
40
Liquids
When heated, the ions move faster and
eventually separate from each other to form a
liquid. The ions are loosely held together by
the oppositely charged ions, but the ions are
moving too fast for the crystal lattice to
stay together.
41
Gases
Since all gas molecules spread out the same way,
equal volumes of gas under equal conditions of
temperature and pressure will contain equal
numbers of molecules of gas. 22.4 L of any gas
at STP (1.00 atm and 273K) will contain one mole
(6.02 X 1023) gas molecules. Since there is
space between gas molecules, gases are affected
by changes in pressure.
42
Types of Matter
  • Substances (Homogeneous)
  • Elements (cannot be decomposed by chemical
    change) Al, Ne, O, Br, H
  • Compounds (can be decomposed by chemical change)
    NaCl, Cu(ClO3)2, KBr, H2O, C2H6
  • Mixtures
  • Homogeneous Solutions (solvent solute)
  • Heterogeneous soil, Italian dressing, etc.

43
Elements
  • A sample of lead atoms (Pb). All atoms in the
    sample consist of lead, so the substance is
    homogeneous.
  • A sample of chlorine atoms (Cl). All atoms in
    the sample consist of chlorine, so the substance
    is homogeneous.

44
Compounds
  • Lead has two charges listed, 2 and 4. This is
    a sample of lead (II) chloride (PbCl2). Two or
    more elements bonded in a whole-number ratio is a
    COMPOUND.
  • This compound is formed from the 4 version of
    lead. This is lead (IV) chloride (PbCl4).
    Notice how both samples of lead compounds have
    consistent composition throughout? Compounds are
    homogeneous!

45
Mixtures
  • A mixture of lead atoms and chlorine atoms. They
    exist in no particular ratio and are not
    chemically combined with each other. They can be
    separated by physical means.
  • A mixture of PbCl2 and PbCl4 formula units.
    Again, they are in no particular ratio to each
    other and can be separated without chemical
    change.

46
Phase Changes
  • Phase Change Types
  • Phase Change Diagrams
  • Heat of Phase Change
  • Evaporation

47
Phase Change Types
48
Phase Change Diagrams
AB Solid Phase BC Melting (S L) CD Liquid
Phase DE Boiling (L G) EF Gas Phase
Notice how temperature remains constant during a
phase change? Thats because the PE is changing,
not the KE.
49
Heat of Phase Change
  • How many joules would it take to melt 100. g of
    H2O (s) at 0oC?
  • qmHf (100. g)(334 J/g) 33400 J
  • How many joules would it take to boil 100. g of
    H2O (l) at 100oC?
  • qmHv (100.g)(2260 J/g) 226000 J

50
Evaporation
  • When the surface molecules of a gas travel
    upwards at a great enough speed to escape.
  • The pressure a vapor exerts when sealed in a
    container at equilibrium is called vapor
    pressure, and can be found on Table H.
  • When the liquid is heated, its vapor pressure
    increases.
  • When the liquids vapor pressure equals the
    pressure exerted on it by the outside atmosphere,
    the liquid can boil.
  • If the pressure exerted on a liquid increases,
    the boiling point of the liquid increases
    (pressure cooker). If the pressure decreases,
    the boiling point of the liquid decreases
    (special cooking directions for high elevations).

51
Reference Table H Vapor Pressure of Four Liquids
52
Bonding
  • 1) The Periodic Table
  • 2) Ions
  • 3) Ionic Bonding
  • 4) Covalent Bonding
  • 5) Metallic Bonding

53
The Periodic Table
  • Metals
  • Nonmetals
  • Metalloids
  • Chemistry of Groups
  • Electronegativity
  • Ionization Energy

54
Metals
  • Have luster, are malleable and ductile, good
    conductors of heat and electricity
  • Lose electrons to nonmetal atoms to form
    positively charged ions in ionic bonds
  • Large atomic radii compared to nonmetal atoms
  • Low electronegativity and ionization energy
  • Left side of the periodic table (except H)

55
Nonmetals
  • Are dull and brittle, poor conductors
  • Gain electrons from metal atoms to form
    negatively charged ions in ionic bonds
  • Share unpaired valence electrons with other
    nonmetal atoms to form covalent bonds and
    molecules
  • Small atomic radii compared to metal atoms
  • High electronegativity and ionization energy
  • Right side of the periodic table (except Group 18)

56
Metalloids
  • Found lying on the jagged line between metals and
    nonmetals flatly touching the line (except Al and
    Po).
  • Share properties of metals and nonmetals (Si is
    shiny like a metal, brittle like a nonmetal and
    is a semiconductor).

57
Chemistry of Groups
  • Group 1 Alkali Metals
  • Group 2 Alkaline Earth Metals
  • Groups 3-11 Transition Elements
  • Group 17 Halogens
  • Group 18 Noble Gases
  • Diatomic Molecules

58
Group 1 Alkali Metals
  • Most active metals, only found in compounds in
    nature
  • React violently with water to form hydrogen gas
    and a strong base 2 Na (s) H2O (l) ? 2 NaOH
    (aq) H2 (g)
  • 1 valence electron
  • Form 1 ion by losing that valence electron
  • Form oxides like Na2O, Li2O, K2O

59
Group 2 Alkaline Earth Metals
  • Very active metals, only found in compounds in
    nature
  • React strongly with water to form hydrogen gas
    and a base
  • Ca (s) 2 H2O (l) ? Ca(OH)2 (aq) H2 (g)
  • 2 valence electrons
  • Form 2 ion by losing those valence electrons
  • Form oxides like CaO, MgO, BaO

60
Groups 3-11 Transition Metals
  • Many can form different possible charges of ions
  • If there is more than one ion listed, give the
    charge as a Roman numeral after the name
  • Cu1 copper (I) Cu2 copper (II)
  • Compounds containing these metals can be colored.

61
Group 17 Halogens
  • Most reactive nonmetals
  • React violently with metal atoms to form halide
    compounds 2 Na Cl2 ? 2 NaCl
  • Only found in compounds in nature
  • Have 7 valence electrons
  • Gain 1 valence electron from a metal to form -1
    ions
  • Share 1 valence electron with another nonmetal
    atom to form one covalent bond.

62
Group 18 Noble Gases
  • Are completely nonreactive since they have eight
    valence electrons, making a stable octet.
  • Kr and Xe can be forced, in the laboratory, to
    give up some valence electrons to react with
    fluorine.
  • Since noble gases do not naturally bond to any
    other elements, one atom of noble gas is
    considered to be a molecule of noble gas. This
    is called a monatomic molecule. Ne represents an
    atom of Ne and a molecule of Ne.

63
Diatomic Molecules
  • Br, I, N, Cl, H, O and F are so reactive that
    they exist in a more chemically stable state when
    they covalently bond with another atom of their
    own element to make two-atom, or diatomic
    molecules.
  • Br2, I2, N2, Cl2, H2, O2 and F2
  • The decomposition of water 2 H2O ? 2 H2 O2

64
Electronegativity
  • An atoms attraction to electrons in a chemical
    bond.
  • F has the highest, at 4.0
  • Fr has the lowest, at 0.7
  • If two atoms that are different in EN (END) from
    each other by 1.7 or more collide and bond (like
    a metal atom and a nonmetal atom), the one with
    the higher electronegativity will pull the
    valence electrons away from the atom with the
    lower electronegativity to form a (-) ion. The
    atom that was stripped of its valence electrons
    forms a () ion.
  • If the two atoms have an END of less than 1.7,
    they will share their unpaired valence
    electrons…covalent bond!

65
Ionization Energy
  • The energy required to remove the most loosely
    held valence electron from an atom in the gas
    phase.
  • High electronegativity means high ionization
    energy because if an atom is more attracted to
    electrons, it will take more energy to remove
    those electrons.
  • Metals have low ionization energy. They lose
    electrons easily to form () charged ions.
  • Nonmetals have high ionization energy but high
    electronegativity. They gain electrons easily to
    form (-) charged ions when reacted with metals,
    or share unpaired valence electrons with other
    nonmetal atoms.

66
Ions
  • Ions are charged particles formed by the gain or
    loss of electrons.
  • Metals lose electrons (oxidation) to form ()
    charged cations.
  • Nonmetals gain electrons (reduction) to form (-)
    charged anions.
  • Atoms will gain or lose electrons in such a way
    that they end up with 8 valence electrons (stable
    octet).
  • The exceptions to this are H, Li, Be and B, which
    are not large enough to support 8 valence
    electrons. They must be satisfied with 2 (Li,
    Be, B) or 0 (H).

67
Metal Ions (Cations)
  • Na 2-8-1
  • Na1 2-8
  • Ca 2-8-8-2
  • Ca2 2-8-8
  • Al 2-8-3
  • Al3 2-8

Note that when the atom loses its valence
electron, the next lower PEL becomes the valence
PEL. Notice how the dot diagrams for metal
ions lack dots! Place brackets around the
element symbol and put the charge on the upper
right outside!
68
Nonmetal Ions (Anions)
Note how the ions all have 8 valence electrons.
Also note the gained electrons as red dots.
Nonmetal ion dot diagrams show 8 dots, with
brackets around the dot diagram and the charge of
the ion written to the upper right side outside
the brackets.
  • F 2-7
  • F-1 2-8
  • O 2-6
  • O-2 2-8
  • N 2-5
  • N-3 2-8

69
Ionic Bonding
  • If two atoms that are different in EN (END) from
    each other by 1.7 or more collide and bond (like
    a metal atom and a nonmetal atom), the one with
    the higher electronegativity will pull the
    valence electrons away from the atom with the
    lower electronegativity to form a (-) ion. The
    atom that was stripped of its valence electrons
    forms a () ion.
  • The oppositely charged ions attract to form the
    bond. It is a surface bond that can be broken by
    melting or dissolving in water.
  • Ionic bonding forms ionic crystal lattices, not
    molecules.

70
Example of Ionic Bonding
71
Covalent Bonding
  • If two nonmetal atoms have an END of 1.7 or less,
    they will share their unpaired valence electrons
    to form a covalent bond.
  • A particle made of covalently bonded nonmetal
    atoms is called a molecule.
  • If the END is between 0 and 0.4, the sharing of
    electrons is equal, so there are no charged ends.
    This is NONPOLAR covalent bonding.
  • If the END is between 0.5 and 1.7, the sharing of
    electrons is unequal. The atom with the higher
    EN will be d- and the one with the lower EN will
    be d charged. This is a POLAR covalent bonding.
    (d means partial)

72
Examples of Covalent Bonding
73
Metallic Bonding
  • Metal atoms of the same element bond with each
    other by sharing valence electrons that they lose
    to each other.
  • This is a lot like an atomic game of hot
    potato, where metal kernals (the atom inside the
    valence electrons) sit in a crystal lattice,
    passing valence electrons back and forth between
    each other).
  • Since electrons can be forced to travel in a
    certain direction within the metal, metals are
    very good at conducting electricity in all phases.

74
Compounds
  • 1) Types of Compounds
  • 2) Formula Writing
  • 3) Formula Naming
  • 4) Empirical Formulas
  • 5) Molecular Formulas
  • 6) Types of Chemical Reactions
  • 7) Balancing Chemical Reactions
  • 8) Attractive Forces

75
Types of Compounds
  • Ionic made of metal and nonmetal ions. Form an
    ionic crystal lattice when in the solid phase.
    Ions separate when melted or dissolved in water,
    allowing electrical conduction. Examples NaCl,
    K2O, CaBr2
  • Molecular made of nonmetal atoms bonded to form
    a distinct particle called a molecule. Bonds do
    not break upon melting or dissolving, so
    molecular substances do not conduct electricity.
    EXCEPTION Acids HA- (aq) ionize in water to
    form H3O and A-, so they do conduct.
  • Network made up of nonmetal atoms bonded in a
    seemingly endless matrix of covalent bonds with
    no distinguishable molecules. Very high m.p.,
    dont conduct.

76
Ionic Compounds
77
Molecular Compounds
78
Network Solids
Network solids are made of nonmetal atoms
covalently bonded together to form large crystal
lattices. No individual molecules can be
distinguished. Examples include C (diamond) and
SiO2 (quartz). Corundum (Al2O3) also forms
these, even though Al is considered a metal.
Network solids are among the hardest materials
known. They have extremely high melting points
and do not conduct electricity.
79
Formula Writing
  • The charge of the () ion and the charge of the
    (-) ion must cancel out to make the formula. Use
    subscripts to indicate how many atoms of each
    element there are in the compound, no subscript
    if there is only one atom of that element.
  • Na1 and Cl-1 NaCl
  • Ca2 and Br-1 CaBr2
  • Al3 and O-2 Al2O3
  • Zn2 and PO4-3 Zn3(PO4)2
  • Try these problems!

80
Formulas to Write
  • Ba2 and N-3
  • NH41 and SO4-2
  • Li1 and S-2
  • Cu2 and NO3-1
  • Al3 and CO3-2
  • Fe3 and Cl-1
  • Pb4 and O-2
  • Pb2 and O-2

81
Formula Naming
  • Compounds are named from the elements or
    polyatomic ions that form them.
  • KCl potassium chloride
  • Na2SO4 sodium sulfate
  • (NH4)2S ammonium sulfide
  • AgNO3 silver nitrate
  • Notice all the metals listed here only have one
    charge listed? So what do you do if a metal has
    more than one charge listed? Take a peek!

82
The Stock System
  • CrCl2 chromium (II) chloride Try
  • CrCl3 chromium (III) chloride Co(NO3)2
    and
  • CrCl6 chromium (VI) chloride Co(NO3)3
  • FeO iron (II) oxide MnS manganese (II)
    sulfide
  • Fe2O3 iron (III) oxide MnS2 manganese (IV)
    sulfide
  • The Roman numeral is the charge of the metal ion!

83
Empirical Formulas
  • Ionic formulas represent the simplest whole
    number mole ratio of elements in a compound.
  • Ca3N2 means a 32 ratio of Ca ions to N ions in
    the compound.
  • Many molecular formulas can be simplified to
    empirical formulas
  • Ethane (C2H6) can be simplified to CH3. This is
    the empirical formula…the ratio of C to H in the
    molecule.
  • All ionic compounds have empirical formulas.

84
Molecular Formulas
  • The count of the actual number of atoms of each
    element in a molecule.
  • H2O a molecule made of two H atoms and one O
    atom covalently bonded together.
  • C2H6O A molecule made of two C atoms, six H
    atoms and one O atom covalently bonded together.
  • Molecular formulas are whole-number multiples of
    empirical formulas
  • H2O 1 X (H2O)
  • C8H16 8 X (CH2)
  • Calculating Molecular Formulas

85
Types of Chemical Reactions
  • Redox Reactions driven by the loss (oxidation)
    and gain (reduction) of electrons. Any species
    that does not change charge is called the
    spectator ion.
  • Synthesis
  • Decomposition
  • Single Replacement
  • Ion Exchange Reaction driven by the formation
    of an insoluble precipitate. The ions that
    remain dissolved throughout are the spectator
    ions.
  • Double Replacement

86
Synthesis
  • Two elements combine to form a compound
  • 2 Na O2 ? Na2O
  • Same reaction, with charges added in
  • 2 Na0 O20 ? Na21O-2
  • Na0 is oxidized (loses electrons), is the
    reducing agent
  • O20 is reduced (gains electrons), is the
    oxidizing agent
  • Electrons are transferred from the Na0 to the
    O20.
  • No spectator ions, there are only two elements
    here.

87
Decomposition
  • A compound breaks down into its original
    elements.
  • Na2O ? 2 Na O2
  • Same reaction, with charges added in
  • Na21O-2 ? 2 Na0 O20
  • O-2 is oxidized (loses electrons), is the
    reducing agent
  • Na1 is reduced (gains electrons), is the
    oxidizing agent
  • Electrons are transferred from the O-2 to the
    Na1.
  • No spectator ions, there are only two elements
    here.

88
Single Replacement
  • An element replaces the same type of element in a
    compound.
  • Ca 2 KCl ? CaCl2 2 K
  • Same reaction, with charges added in
  • Ca0 2 K1Cl-1 ? Ca2Cl2-1 2 K0
  • Ca0 is oxidized (loses electrons), is the
    reducing agent
  • K1 is reduced (gains electrons), is the
    oxidizing agent
  • Electrons are transferred from the Ca0 to the
    K1.
  • Cl-1 is the spectator ion, since its charge
    doesnt change.

89
Double Replacement
  • The () ion of one compound bonds to the (-) ion
    of another compound to make an insoluble
    precipitate. The compounds must both be
    dissolved in water to break the ionic bonds
    first.
  • NaCl (aq) AgNO3 (aq) ? NaNO3 (aq) AgCl (s)
  • The Cl-1 and Ag1 come together to make the
    insoluble precipitate, which looks like snow in
    the test tube.
  • No species change charge, so this is not a redox
    reaction.
  • Since the Na1 and NO3-1 ions remain dissolved
    throughout the reaction, they are the spectator
    ions.
  • How do identify the precipitate?

90
Identifying the Precipitate
  • The precipitate is the compound that is
    insoluble. AgCl is a precipitate because Cl- is
    a halide. Halides are soluble, except when
    combined with Ag and others.

91
Balancing Chemical Reactions
  • Balance one element or ion at a time
  • Use a pencil
  • Use coefficients only, never change formulas
  • Revise if necessary
  • The coefficient multiplies everything in the
    formula by that amount
  • 2 Ca(NO3)2 means that you have 2 Ca, 4 N and 12
    O.
  • Examples for you to try!

92
Reactions to Balance
  • ___NaCl ? ___Na ___Cl2
  • ___Al ___O2 ? ___Al2O3
  • ___SO3 ? ___SO2 ___O2
  • ___Ca ___HNO3 ? ___Ca(NO3)2 ___H2
  • __FeCl3 __Pb(NO3)2 ? __Fe(NO3)3 __PbCl2

93
Attractive Forces
  • Molecules have partially charged ends. The d
    end of one molecule attracts to the d- end of
    another molecule.
  • Ions are charged () or (-). Positively charged
    ions attract other to form ionic bonds, a type of
    attractive force.
  • Since partially charged ends result in weaker
    attractions than fully charged ends, ionic
    compounds generally have much higher melting
    points than molecular compounds.
  • Determining Polarity of Molecules
  • Hydrogen Bond Attractions

94
Determining Polarity of Molecules
--------------------------------------------------
---------------------------
95
Hydrogen Bond Attractions
A hydrogen bond attraction is a very strong
attractive force between the H end of one polar
molecule and the N, O or F end of another polar
molecule. This attraction is so strong that
water is a liquid at a temperature where most
compounds that are much heavier than water (like
propane, C3H8) are gases. This also gives water
its surface tension and its ability to form a
meniscus in a narrow glass tube.
96
Math of Chemistry
  • 1) Formula Mass
  • 2) Percent Composition
  • 3) Mole Problems
  • 4) Gas Laws
  • 5) Neutralization
  • 6) Concentration
  • 7) Significant Figures and Rounding
  • 8) Metric Conversions
  • 9) Calorimetry

97
Formula Mass
  • Gram Formula Mass sum of atomic masses of all
    elements in the compound
  • Round given atomic masses to the nearest tenth
  • H2O (2 X 1.0) (1 X 16.0) 18.0 grams/mole
  • Na2SO4 (2 X 23.0)(1 X 32.1)(4 X 16.0) 142.1
    g/mole
  • Now you try
  • BaBr2
  • CaSO4
  • Al2(CO3)3

98
Percent Composition
The mass of part is the number of atoms of that
element in the compound. The mass of whole is
the formula mass of the compound. Dont forget
to take atomic mass to the nearest tenth! This
is a problem for you to try.
99
Practice Percent Composition Problem
  • What is the percent by mass of each element in
    Li2SO4?

100
Mole Problems
  • Grams Moles
  • Molecular Formula
  • Stoichiometry

101
Grams Moles
  • How many grams will 3.00 moles of NaOH (40.0
    g/mol) weigh?
  • 3.00 moles X 40.0 g/mol 120. g
  • How many moles of NaOH (40.0 g/mol) are
    represented by 10.0 grams?
  • (10.0 g) / (40.0 g/mol) 0.250 mol

102
Molecular Formula
  • Molecular Formula (Molecular Mass/Empirical
    Mass) X Empirical Formula
  • What is the molecular formula of a compound with
    an empirical formula of CH2 and a molecular mass
    of 70.0 grams/mole?
  • 1) Find the Empirical Formula Mass CH2 14.0
  • 2) Divide the MM/EM 70.0/14.0 5
  • 3) Multiply the molecular formula by the result
  • 5 (CH2) C5H10

103
Stoichiometry
  • Moles of Target Moles of Given X (Coefficent of
    Target/Coefficient of given)
  • Given the balanced equation N2 3 H2 ? 2 NH3,
    How many moles of H2 need to be completely
    reacted with N2 to yield 20.0 moles of NH3?
  • 20.0 moles NH3 X (3 H2 / 2 NH3) 30.0 moles H2

104
Gas Laws
  • Make a data table to put the numbers so you can
    eliminate the words.
  • Make sure that any Celsius temperatures are
    converted to Kelvin (add 273).
  • Rearrange the equation before substituting in
    numbers. If you are trying to solve for T2, get
    it out of the denominator first by
    cross-multiplying.
  • If one of the variables is constant, then
    eliminate it.
  • Try these problems!

105
Gas Law Problem 1
  • A 2.00 L sample of N2 gas at STP is compressed to
    4.00 atm at constant temp-erature. What is the
    new volume of the gas?
  • V2 P1V1 / P2
  • (1.00 atm)(2.00 L) / (4.00 atm)
  • 0.500 L

106
Gas Law Problem 2
  • To what temperature must a 3.000 L sample of O2
    gas at 300.0 K be heated to raise the volume to
    10.00 L?
  • T2 V2T1/V1
  • (10.00 L)(300.0 K) / (3.000 L) 1000. K

107
Gas Law Problem 3
  • A 3.00 L sample of NH3 gas at 100.0 kPa is cooled
    from 500.0 K to 300.0 K and its pressure is
    reduced to 80.0 kPa. What is the new volume of
    the gas?
  • V2 P1V1T2 / P2T1
  • (100.0 kPa)(3.00 L)(300. K) / (80.0 kPa)(500.
    K)
  • 2.25 L

108
Neutralization
  • 10.0 mL of 0.20 M HCl is neutralized by 40.0 mL
    of NaOH. What is the concentration of the NaOH?
  • H MaVa OH MbVb, so Mb H MaVa / OH Vb
  • (1)(0.20 M)(10.0 mL) / (1) (40.0 mL) 0.050 M
  • How many mL of 2.00 M H2SO4 are needed to
    completely neutralize 30.0 mL of 0.500 M KOH?

109
Concentration
  • Molarity
  • Parts per Million
  • Percent by Mass
  • Percent by Volume

110
Molarity
  • What is the molarity of a 500.0 mL solution of
    NaOH (FM 40.0) with 60.0 g of NaOH (aq)?
  • Convert g to moles and mL to L first!
  • M moles / L 1.50 moles / 0.5000 L 3.00 M
  • How many grams of NaOH does it take to make 2.0 L
    of a 0.100 M solution of NaOH (aq)?
  • Moles M X L 0.100 M X 2.0 L 0.200 moles
  • Convert moles to grams 0.200 moles X 40.0 g/mol
    8.00 g

111
Parts Per Million
  • 100.0 grams of water is evaporated and analyzed
    for lead. 0.00010 grams of lead ions are found.
    What is the concentration of the lead, in parts
    per million?
  • ppm (0.00010 g) / (100.0 g) X 1 000 000 1.0
    ppm
  • If the legal limit for lead in the water is 3.0
    ppm, then the water sample is within the legal
    limits (its OK!)

112
Percent by Mass
  • A 50.0 gram sample of a solution is evaporated
    and found to contain 0.100 grams of sodium
    chloride. What is the percent by mass of sodium
    chloride in the solution?
  • Comp (0.100 g) / (50.0 g) X 100 0.200

113
Percent By Volume
  • Substitute volume for mass in the above
    equation.
  • What is the percent by volume of hexane if 20.0
    mL of hexane are dissolved in benzene to a total
    volume of 80.0 mL?
  • Comp (20.0 mL) / (80.0 mL) X100 25.0

114
Sig Figs and Rounding
  • How many Significant Figures does a number have?
  • What is the precision of my measurement?
  • How do I round off answers to addition and
    subtraction problems?
  • How do I round off answers to multiplication and
    division problems?

115
How many Sig Figs?
  • Start counting sig figs at the first non-zero.
  • All digits except place-holding zeroes are sig
    figs.

116
What Precision?
  • A numbers precision is determined by the
    furthest (smallest) place the number is recorded
    to.
  • 6000 mL thousands place
  • 6000. mL ones place
  • 6000.0 mL tenths place
  • 5.30 mL hundredths place
  • 8.7 mL tenths place
  • 23.740 mL thousandths place

117
Rounding with addition and subtraction
  • Answers are rounded to the least precise place.

118
Rounding with multiplication and division
  • Answers are rounded to the fewest number of
    significant figures.

119
Metric Conversions
  • Determine how many powers of ten difference there
    are between the two units (no prefix 100) and
    create a conversion factor. Multiply or divide
    the given by the conversion factor.

How many kg are in 38.2 cg? (38.2 cg) /(100000
cg/kg) 0.000382 km How many mL in 0.988
dL? (0.988 dg) X (100 mL/dL) 98.8 mL
120
Calorimetry
  • This equation can be used to determine any of the
    variables here. You will not have to solve for
    C, since we will always assume that the energy
    transfer is being absorbed by or released by a
    measured quantity of water, whose specific heat
    is given above.
  • Solving for q
  • Solving for m
  • Solving for DT

121
Solving for q
  • How many joules are absorbed by 100.0 grams of
    water in a calorimeter if the temperature of the
    water increases from 20.0oC to 50.0oC?
  • q mCDT (100.0 g)(4.18 J/goC)(30.0oC) 12500 J

122
Solving for m
  • A sample of water in a calorimeter cup increases
    from 25oC to 50.oC by the addition of 500.0
    joules of energy. What is the mass of water in
    the calorimeter cup?
  • q mCDT, so m q / CDT (500.0 J) / (4.18
    J/goC)(25oC) 4.8 g

123
Solving for DT
  • If a 50.0 gram sample of water in a calorimeter
    cup absorbs 1000.0 joules of energy, how much
    will the temperature rise by?
  • q mCDT, so DT q / mC (1000.0 J)/(50.0
    g)(4.18 J/goC) 4.8oC
  • If the water started at 20.0oC, what will the
    final temperature be?
  • Since the water ABSORBS the energy, its
    temperature will INCREASE by the DT 20.0oC
    4.8oC 24.8oC

124
Kinetics and Thermodynamics
  • 1) Reaction Rate
  • 2) Heat of Reaction
  • 3) Potential Energy Diagrams
  • 4) Equilibrium
  • 5) Le Châteliers Principle
  • 6) Solubility Curves

125
Reaction Rate
  • Reactions happen when reacting particles collide
    with sufficient energy (activation energy) and at
    the proper angle.
  • Anything that makes more collisions in a given
    time will make the reaction rate increase.
  • Increasing temperature
  • Increasing concentration (pressure for gases)
  • Increasing surface area (solids)
  • Adding a catalyst makes a reaction go faster by
    removing steps from the mechanism and lowering
    the activation energy without getting used up in
    the process.

126
Heat of Reaction
  • Reactions either absorb PE (endothermic, DH) or
    release PE (exothermic, -DH)

Exothermic, PE?KE, Temp? Endothermic, KE?PE, Temp?
Rewriting the equation with heat included 4
Al(s) 3 O2(g) ? 2 Al2O3(s) 3351 kJ N2(g)
O2(g) 182.6 kJ ? 2 NO(g)
127
Potential Energy Diagrams
  • Steps of a reactions
  • Reactants have a certain amount of PE stored in
    their bonds (Heat of Reactants)
  • The reactants are given enough energy to collide
    and react (Activation Energy)
  • The resulting intermediate has the highest energy
    that the reaction can make (Heat of Activated
    Complex)
  • The activated complex breaks down and forms the
    products, which have a certain amount of PE
    stored in their bonds (Heat of Products)
  • Hproducts - Hreactants DH
    EXAMPLES

128
Making a PE Diagram
  • X axis Reaction Coordinate (time, no units)
  • Y axis PE (kJ)
  • Three lines representing energy (Hreactants,
    Hactivated complex, Hproducts)
  • Two arrows representing energy changes
  • From Hreactants to Hactivated complex Activation
    Energy
  • From Hreactants to Hproducts DH
  • ENDOTHERMIC PE DIAGRAM
  • EXOTHERMIC PE DIAGRAM

129
Endothermic PE Diagram
If a catalyst is added?
130
Endothermic with Catalyst
The red line represents the catalyzed reaction.
131
Exothermic PE Diagram
What does it look like with a catalyst?
132
Exothermic with a Catalyst
The red line represents the catalyzed reaction.
Lower A.E. and faster reaction time!
133
Equilibrium
When the rate of the forward reaction equals the
rate of the reverse reaction.
134
Examples of Equilibrium
  • Solution Equilibrium when a solution is
    saturated, the rate of dissolving equals the rate
    of precipitating.
  • NaCl (s) ? Na1 (aq) Cl-1 (aq)
  • Vapor-Liquid Equilibrium when a liquid is
    trapped with air in a container, the liquid
    evaporates until the rate of evaporation equals
    the rate of condensation.
  • H2O (l) ? H2O (g)
  • Phase equilibrium At the melting point, the
    rate of solid turning to liquid equals the rate
    of liquid turning back to solid.
  • H2O (s) ? H2O (l)

135
Le Châteliers Principle
  • If a system at equilibrium is stressed, the
    equilibrium will shift in a direction that
    relieves that stress.
  • A stress is a factor that affects reaction rate.
    Since catalysts affect both reaction rates
    equally, catalysts have no effect on a system
    already at equilibrium.
  • Equilibrium will shift AWAY from what is added
  • Equilibrium will shift TOWARDS what is removed.
  • This is because the shift will even out the
    change in reaction rate and bring the system back
    to equilibrium
  • NEXT

136
Steps to Relieving Stress
  • 1) Equilibrium is subjected to a STRESS.
  • 2) System SHIFTS towards what is removed from the
    system or away from what is added.
  • The shift results in a CHANGE OF CONCENTRATION
    for both the products and the reactants.
  • If the shift is towards the products, the
    concentration of the products will increase and
    the concentration of the reactants will decrease.
  • If the shift is towards the reactants, the
    concentration of the reactants will increase and
    the concentration of the products will decrease.
  • NEXT

137
Examples
  • For the reaction N2(g) 3H2(g) ? 2 NH3(g) heat
  • Adding N2 will cause the equilibrium to shift
    RIGHT, resulting in an increase in the
    concentration of NH3 and a decrease in the
    concentration of N2 and H2.
  • Removing H2 will cause a shift to the LEFT,
    resulting in a decrease in the concentration of
    NH3 and an increase in the concentration of N2
    and H2.
  • Increasing the temperature will cause a shift to
    the LEFT, same results as the one above.
  • Decreasing the pressure will cause a shift to the
    LEFT, because there is more gas on the left side,
    and making more gas will bring the pressure back
    up to its equilibrium amount.
  • Adding a catalyst will have no effect, so no
    shift will happen.

138
Solubility Curves
  • Solubility the maximum quantity of solute that
    can be dissolved in a given quantity of solvent
    at a given temperature to make a saturated
    solution.
  • Saturated a solution containing the maximum
    quantity of solute that the solvent can hold.
    The limit of solubility.
  • Supersaturated the solution is holding more
    than it can theoretically hold OR there is excess
    solute which precipitates out. True
    supersaturation is rare.
  • Unsaturated There are still solvent molecules
    available to dissolve more solute, so more can
    dissolve.
  • How ionic solutes dissolve in water polar water
    molecules attach to the ions and tear them off
    the crystal.

139
Solubility
Solubility go to the temperature and up to the
desired line, then across to the Y-axis. This is
how many g of solute are needed to make a
saturated solution of that solute in 100g of H2O
at that particular temperature. At 40oC, the
solubility of KNO3 in 100g of water is 64 g. In
200g of water, double that amount. In 50g of
water, cut it in half.
140
Supersaturated
If 120 g of NaNO3 are added to 100g of water at
30oC 1) The solution would be SUPERSATURATED,
because there is more solute dissolved than the
solubility allows 2) The extra 25g would
precipitate out 3) If you heated the solution up
by 24oC (to 54oC), the excess solute would
dissolve.
141
Unsaturated
If 80 g of KNO3 are added to 100g of water at
60oC 1) The solution would be UNSATURATED,
because there is less solute dissolved than the
solubility allows 2) 26g more can be added to
make a saturated solution 3) If you cooled the
solution down by 12oC (to 48oC), the solution
would become saturated
142
How Ionic Solutes Dissolve in Water
Water solvent molecules attach to the ions (H end
to the Cl-, O end to the Na)
Water solvent holds the ions apart and keeps the
ions from coming back together
143
Acids and Bases
  • 1) Formulas, Naming and Properties of Acids
  • 2) Formulas, Naming and Properties of Bases
  • 3) Neutralization
  • 4) pH
  • 5) Indicators
  • 6) Alternate Theories

144
Formulas, Naming and Properties of Acids
  • Arrhenius Definition of Acids molecules that
    dissolve in water to produce H3O (hydronium) as
    the only positively charged ion in solution.
  • HCl (g) H2O (l) ? H3O (aq) Cl-
  • Properties of Acids
  • Naming of Acids
  • Formula Writing of Acids

145
Properties of Acids
  • Acids react with metals above H2 on Table J to
    form H2(g) and a salt.
  • Acids have a pH of less than 7.
  • Dilute solutions of acids taste sour.
  • Acids turn phenolphthalein CLEAR, litmus RED and
    bromthymol blue YELLOW.
  • Acids neutralize bases.
  • Acids are formed when acid anhydrides (NO2, SO2,
    CO2) react with water for form acids. This is
    how acid rain forms from auto and industrial
    emissions.

146
Naming of Acids
  • Binary Acids (H and a nonmetal)
  • hydro (nonmetal) -ide ic acid
  • HCl (aq) hydrochloric acid
  • Ternary Acids (H and a polyatomic ion)
  • (polyatomic ion) -ate ic acid
  • HNO3 (aq) nitric acid
  • (polyatomic ion) -ide ic acid
  • HCN (aq) cyanic acid
  • (polyatomic ion) -ite ous acid
  • HNO2 (aq) nitrous acid

147
Formula Writing of Acids
  • Acids formulas get written like any other. Write
    the H1 first, then figure out what the negative
    ion is based on the name. Cancel out the charges
    to write the formula. Dont forget the (aq)
    after it…its only an acid if its in water!
  • Hydrosulfuric acid H1 and S-2 H2S (aq)
  • Carbonic acid H1 and CO3-2 H2CO3 (aq)
  • Chlorous acid H1 and ClO2-1 HClO2 (aq)
  • Hydrobromic acid H1 and Br-1 HBr (aq)
  • Hydronitric acid
  • Hypochlorous acid
  • Perchloric acid

148
Formulas, Naming and Properties of Bases
  • Arrhenius Definition of Bases ionic compounds
    that dissolve in water to produce OH- (hydroxide)
    as the only negatively charged ion in solution.
  • NaOH (s) ? Na1 (aq) OH-1 (aq)
  • Properties of Bases
  • Naming of Bases
  • Formula Writing of Bases

149
Properties of Bases
  • Bases react with fats to form soap and glycerol.
    This process is called saponification.
  • Bases have a pH of more than 7.
  • Dilute solutions of bases taste bitter.
  • Bases turn phenolphthalein PINK, litmus BLUE and
    bromthymol blue BLUE.
  • Bases neutralize acids.
  • Bases are formed when alkali metals or alkaline
    earth metals react with water. The words
    alkali and alkaline mean basic, as opposed
    to acidic.

150
Naming of Bases
  • Bases are named like any ionic compound, the name
    of the metal ion first (with a Roman numeral if
    necessary) followed by hydroxide.

Fe(OH)2 (aq) iron (II) hydroxide Fe(OH)3 (aq)
iron (III) hydroxide Al(OH)3 (aq) aluminum
hydroxide NH3 (aq) is the same thing as
NH4OH NH3 H2O ? NH4OH Also called ammonium
hydroxide.
151
Formula Writing of Bases
  • Formula writing of bases is the same as for any
    ionic formula writing. The charges of the ions
    have to cancel out.
  • Calcium hydroxide Ca2 and OH-1 Ca(OH)2 (aq)
  • Potassium hydroxide K1 and OH-1 KOH (aq)
  • Lead (II) hydroxide Pb2 and OH-1 Pb(OH)2
    (aq)
  • Lead (IV) hydroxide Pb4 and OH-1 Pb(OH)4
    (aq)
  • Lithium hydroxide
  • Copper (II) hydroxide
  • Magnesium hydroxide

152
Neutralization
  • H1 OH-1 ? HOH
  • Acid Base ? Water Salt (double replacement)
  • HCl (aq) NaOH (aq) ? HOH (l) NaCl (aq)
  • H2SO4 (aq) KOH (aq) ? 2 HOH (l) K2SO4 (aq)
  • HBr (aq) LiOH (aq) ?
  • H2CrO4 (aq) NaOH (aq) ?
  • HNO3 (aq) Ca(OH)2 (aq) ?
  • H3PO4 (aq) Mg(OH)2 (aq) ?

153
pH
  • A change of 1 in pH is a tenfold increase in acid
    or base strength.
  • A pH of 4 is 10 times more acidic than a pH of 5.
  • A pH of 12 is 100 times more basic than a pH of
    10.

154
Indicators
At a pH of 2 Methyl Orange red Bromthymol Blue
yellow Phenolphthalein colorless Litmus
red Bromcresol Green yellow Thymol Blue yellow
Methyl orange is red at a pH of 3.2 and below
and yellow at a pH of 4.4 and higher. In between
the two numbers, it is an intermediate color that
is not listed on this table.
155
Alternate Theories
  • Arrhenius Theory acids and bases must be in
    aqueous solution.
  • Alternate Theory Not necessarily so!
  • Acid proton (H1) donor…gives up H1 in a
    reaction.
  • Base proton (H1) acceptor…gains H1 in a
    reaction.
  • HNO3 H2O ? H3O1 NO3-1
  • Since HNO3 lost an H1 during the reaction, it is
    an acid.
  • Since H2O gained the H1 that HNO3 lost, it is a
    base.

156
Oxidation and Reduction
  • 1) Oxidation Numbers
  • 2) Identifying OX, RD and SI Species
  • 3) Agents
  • 4) Writing Half-Reactions
  • 5) Balancing Half-Reactions
  • 6) Activity Series
  • 7) Voltaic Cells
  • 8) Electrolytic Cells
  • 9) Electroplating

157
Oxidation Numbers
  • Elements have no charge until they bond to other
    elements.
  • Na0, Li0, H20. S0, N20, C600
  • The formula of a compound is such that the
    charges of the elements making up the compound
    all add up to zero.
  • The symbol and charge of an element or polyatomic
    ion is called a SPECIES.
  • Determine the charge of each species in the
    following compounds
  • NaCl KNO3 CuSO4 Fe2(CO3)3

158
Identifying OX, RD, SI Species
  • Ca0 2 H1Cl-1 ? Ca2Cl-12 H20
  • Oxidation loss of electrons. The species
    becomes more positive in charge. For example,
    Ca0 ? Ca2, so Ca0 is the species that is
    oxidized.
  • Reduction gain of electrons. The species
    becomes more negative in charge. For example,
    H1 ? H0, so the H1 is the species that is
    reduced.
  • Spectator Ion no change in charge. The species
    does not gain or lose any electrons. For
    example, Cl-1 ? Cl-1, so the Cl-1 is the
    spectator ion.

159
Agents
  • Ca0 2 H1Cl-1 ? Ca2Cl-12 H20
  • Since Ca0 is being oxidized and H1 is being
    reduced, the electrons must be going from the Ca0
    to the H1.
  • Since Ca0 would not lose electrons (be oxidized)
    if H1 werent there to gain them, H1 is the
    cause, or agent, of Ca0s oxidation. H1 is the
    oxidizing agent.
  • Since H1 would not gain electrons (be reduced)
    if Ca0 werent there to lose them, Ca0 is the
    cause, or agent, of H1s reduction. Ca0 is the
    reducing agent.

160
Writing Half-Reactions
  • Ca0 2 H1Cl-1 ? Ca2Cl-12 H20
  • Oxidation Ca0 ? Ca2 2e-
  • Reduction 2H1 2e- ? H20

The two electrons lost by Ca0 are gained by the
two H1 (each H1 picks up an electron).
PRACTICE SOME!
161
Practice Half-Reactions
  • Dont forget to determine the charge of each
    species first!
  • 4 Li O2 ? 2 Li2O
  • Oxidation Half-Reaction
  • Reduction Half-Reaction
  • Zn Na2SO4 ? ZnSO4 2 Na
  • Oxidation Half-Reaction
  • Reduction Half-Reaction

162
Balancing Half-Reactions
  • Ca0 Fe3 ? Ca2 Fe0
  • Cas charge changes by 2, so double the Fe.
  • Fes charge changes by 3, so triple the Ca.
  • 3 Ca0 2 Fe3 ? 3 Ca2 2 Fe0
  • Try these
  • __Na0 __H1 ? __Na1 __H20
  • (hint balance the H and H2 first!)
  • __Al0 __Cu2 ? __Al3 __Cu0

163
Activity Series
  • For metals, the higher up the chart the element
    is, the more likely it is to be oxidized. This
    is because metals like to lose electrons, and the
    more active a metallic element is, the more
    easily it can lose them.
  • For nonmetals, the higher up the chart the
    element is, the more likely it is to be reduced.
    This is because nonmetals like to gain electrons,
    and the more active a nonmetallic element is, the
    more easily it can gain them.

164
Metal Activity
3 K0 Fe3Cl-13 REACTION
  • Metallic elements start out with a charge of
    ZERO, so they can only be oxidized to form ()
    ions.
  • The higher of two metals MUST undergo oxidation
    in the reaction, or no reaction will happen.
  • The reaction 3 K FeCl3 ? 3 KCl Fe WILL
    happen, because K is being oxidized, and that is
    what Table J says should happen.
  • The reaction Fe 3 KCl ? FeCl3 3 K will NOT
    happen.

Fe0 3 K1Cl-1 NO REACTION
165
Voltaic Cells
  • Produce electrical current using a spontaneous
    redox reaction
  • Used to make batteries!
  • Materials needed two beakers, piece of the
    oxidized metal (anode, - electrode), solution of
    the oxidized metal, piece of the reduced metal
    (cathode, electrode), solution of the reduced
    metal, porous material (salt bridge), solution of
    a salt that does not contain either metal in the
    reaction, wire and a load to make use of the
    generated current!
  • Use Reference Table J to determine the metals to
    use
  • Higher (-) anode Lower () cathode

166
Making Voltaic Cells
More Info!!!
Create Your Own Cell!!!!
167
How It Works
Since Zn is listed above Cu, Zn0 will be oxidized
when it reacts with Cu2. The reaction Zn
CuSO4 ? ZnSO4 Cu
  • The Zn0 anode loses 2 e-, which go up the wire
    and through the load. The Zn0 electrode gets
    smaller as the Zn0 becomes Zn2 and dissolves
    into solution. The e- go into the Cu0, where
    they sit on the outside surface of the Cu0
    cathode and wait for Cu2 from the solution to
    come over so that the e- can jump on to the Cu2
    and reduce it to Cu0. The size of the Cu0
    electrode increases. The negative ions in
    solution go over the salt bridge to the anode
    side to complete the circuit.

168
You Start At The Anode
169
Make Your Own Cell!!!
170
Electrolytic Cells
  • Use electricity to force a nonspontaneous redox
    reaction to take place.
  • Uses for Electrolytic Cells
  • Decomposition of Alkali Metal Compounds
  • Decomposition of Water into Hydrogen and Oxygen
  • Electroplating
  • Differences between Voltaic and Electrolytic
    Cells
  • ANODE Voltaic (-) Electrolytic ()
  • CATHODE Voltaic () Electrolytic (-)
  • Voltaic 2 half-cells, a salt bridge and a load
  • Electrolytic 1 cell, no salt bridge, IS the load

171
Decomposing Alkali Metal Compounds
2 NaCl ? 2 Na Cl2 The Na1 is reduced at the
(-) cathode, picking up an e- from the
battery The Cl-1 is oxidized at the () anode,
the e- being pulled off by the battery (DC)
172
Decomposing Water
2 H2O ? 2 H2 O2 The H is reduced at the (-)
cathode, yielding H2 (g), which is trapped in the
tube. The O-2 is oxidized at the () anode,
yielding O2 (g), which is trapped in the tube.
173
Electroplating
The Ag0 is oxidized to Ag1 when the () end of
the battery strips its electrons off. The Ag1
migrates through the solution towards the (-)
charged cathode (ring), where it picks up an
electron from the battery and forms Ag0, which
coats on to the ring.
174
Organic Chemistry
  • 1) H
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