Title: Mark Rosengartens Amazing Chemistry Powerpoint Presentation
1Mark 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!
2Outline 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)
3The Atom
- 1) Nucleons
- 2) Isotopes
- 3) Natural Radioactivity
- 4) Half-Life
- 5) Nuclear Power
- 6) Electron Configuation
- 7) Development of the Atomic Model
4Nucleons
- 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.
5Isotopes
- 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.
6Weight-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)
7Most 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.
8Natural 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.
9Alpha Decay
- The nucleus ejects two protons and two neutrons.
The atomic mass decreases by 4, the atomic number
decreases by 2. - 23892U ?
10Beta 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 ?
11Positron 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 ?
12Gamma 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.
13Charges of Decay Particles
14Half-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
15Going 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
16Going 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
17Radioactive 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
18Nuclear Power
- Artificial Transmutation
- Particle Accelerators
- Nuclear Fission
- Nuclear Fusion
19Artificial 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.
20Particle 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.
21Nuclear 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.
22Chain Reaction
23Nuclear 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.
24Electron Configuration
- Basic Configuration
- Valence Electrons
- Electron-Dot (Lewis Dot) Diagrams
- Excited vs. Ground State
- What is Light?
25Basic 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)
26Valence 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.
27Electron-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
28Example Dot Diagrams
Carbon can also have this dot diagram, which
it has when it forms organic compounds.
29Excited 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
30What 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
31EXAMPLE 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.
32Development of the Atomic Model
- Thompson Model
- Rutherford Gold Foil Experiment and Model
- Bohr Model
- Quantum-Mechanical Model
33Thompson Model
- The atom is a positively charged diffuse mass
with negatively charged electrons stuck in it.
34Rutherford 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).
35Bohr Model
- Electrons orbit around the nucleus in energy
levels (shells). Atomic bright-line spectra was
the clue.
36Quantum-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.
37Matter
- 1) Properties of Phases
- 2) Types of Matter
- 3) Phase Changes
38Properties 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)
39Solids
The positive and negative ions alternate in the
ionic crystal lattice of NaCl.
40Liquids
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.
41Gases
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.
42Types 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.
43Elements
- 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.
44Compounds
- 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!
45Mixtures
- 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.
46Phase Changes
- Phase Change Types
- Phase Change Diagrams
- Heat of Phase Change
- Evaporation
47Phase Change Types
48Phase 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.
49Heat 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
50Evaporation
- 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).
51Reference Table H Vapor Pressure of Four Liquids
52Bonding
- 1) The Periodic Table
- 2) Ions
- 3) Ionic Bonding
- 4) Covalent Bonding
- 5) Metallic Bonding
53The Periodic Table
- Metals
- Nonmetals
- Metalloids
- Chemistry of Groups
- Electronegativity
- Ionization Energy
54Metals
- 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)
55Nonmetals
- 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)
56Metalloids
- 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).
57Chemistry 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
58Group 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
59Group 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
60Groups 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.
61Group 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.
62Group 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.
63Diatomic 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
64Electronegativity
- 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
electronscovalent bond!
65Ionization 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.
66Ions
- 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).
67Metal 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!
68Nonmetal 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
69Ionic 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.
70Example of Ionic Bonding
71Covalent 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)
72Examples of Covalent Bonding
73Metallic 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.
74Compounds
- 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
75Types 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.
76Ionic Compounds
77Molecular Compounds
78Network 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.
79Formula 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!
80Formulas 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
81Formula 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!
82The 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!
83Empirical 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 formulathe ratio of C to H in the
molecule. - All ionic compounds have empirical formulas.
84Molecular 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
85Types 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
86Synthesis
- 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.
87Decomposition
- 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.
88Single 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.
89Double 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?
90Identifying 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.
91Balancing 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!
92Reactions 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
93Attractive 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
94Determining Polarity ofMolecules
--------------------------------------------------
---------------------------
95Hydrogen BondAttractions
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.
96Math 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
97Formula 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
98Percent 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.
99Practice PercentComposition Problem
- What is the percent by mass of each element in
Li2SO4?
100Mole Problems
- Grams Moles
- Molecular Formula
- Stoichiometry
101Grams 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
102Molecular 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
103Stoichiometry
- 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
104Gas 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!
105Gas 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
106Gas 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
107Gas 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
108Neutralization
- 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?
109Concentration
- Molarity
- Parts per Million
- Percent by Mass
- Percent by Volume
110Molarity
- 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
111Parts 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!)
112Percent 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
113Percent 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
114Sig 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?
115How many Sig Figs?
- Start counting sig figs at the first non-zero.
- All digits except place-holding zeroes are sig
figs.
116What 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
117Rounding with addition and subtraction
- Answers are rounded to the least precise place.
118Rounding with multiplicationand division
- Answers are rounded to the fewest number of
significant figures.
119Metric 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
120Calorimetry
- 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
121Solving 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
122Solving 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
123Solving 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
124Kinetics and Thermodynamics
- 1) Reaction Rate
- 2) Heat of Reaction
- 3) Potential Energy Diagrams
- 4) Equilibrium
- 5) Le Châteliers Principle
- 6) Solubility Curves
125Reaction 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.
126Heat 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)
127Potential 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
128Making 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
129Endothermic PE Diagram
If a catalyst is added?
130Endothermic with Catalyst
The red line represents the catalyzed reaction.
131Exothermic PE Diagram
What does it look like with a catalyst?
132Exothermic with a Catalyst
The red line represents the catalyzed reaction.
Lower A.E. and faster reaction time!
133Equilibrium
When the rate of the forward reaction equals the
rate of the reverse reaction.
134Examples 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)
135Le 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
136Steps 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
137Examples
- 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.
138Solubility 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.
139Solubility
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.
140Supersaturated
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.
141Unsaturated
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
142How 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
143Acids 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
144Formulas, 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
145Properties 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.
146Naming 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
147Formula 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 itits 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
148Formulas, 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
149Properties 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.
150Naming 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.
151Formula 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
152Neutralization
- 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) ?
153pH
- 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.
154Indicators
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.
155Alternate Theories
- Arrhenius Theory acids and bases must be in
aqueous solution. - Alternate Theory Not necessarily so!
- Acid proton (H1) donorgives up H1 in a
reaction. - Base proton (H1) acceptorgains 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.
156Oxidation 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
157Oxidation 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
158Identifying 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.
159Agents
- 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.
160Writing 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!
161Practice 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
162Balancing 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
163Activity 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.
164Metal 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
165Voltaic 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
166Making Voltaic Cells
More Info!!!
Create Your Own Cell!!!!
167How 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.
168You Start At The Anode
169Make Your Own Cell!!!
170Electrolytic 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
171Decomposing AlkaliMetal 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)
172Decomposing 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.
173Electroplating
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.
174Organic Chemistry