Chemistry Midterm Review Presentation!

1
Chemistry Midterm Review 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!

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.)

3
The Atom

• 1) Nucleons
• 2) Isotopes
• 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
Electron Configuration

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

9
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)

10
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.

11
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

12
Example Dot Diagrams
Carbon can also have this dot diagram, which
it has when it forms organic compounds.
13
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

14
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

15
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.
16
Development of the Atomic Model

• Thompson Model
• Rutherford Gold Foil Experiment and Model
• Bohr Model
• Quantum-Mechanical Model

17
Thompson Model

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

18
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).
19
Bohr Model

• Electrons orbit around the nucleus in energy
  levels (shells). Atomic bright-line spectra was
  the clue.

20
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.

21
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)

22
Solids
The positive and negative ions alternate in the
ionic crystal lattice of NaCl.
23
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.
24
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.
25
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.

26
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.

27
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!

28
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.

29
Bonding

• 1) The Periodic Table
• 2) Ions
• 3) Ionic Bonding
• 4) Covalent Bonding
• 5) Metallic Bonding

30
The Periodic Table

• Metals
• Nonmetals
• Metalloids
• Chemistry of Groups
• Electronegativity
• Ionization Energy

31
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)

32
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)

33
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).

34
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

35
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

36
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

37
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.

38
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.

39
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.

40
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

41
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
  electronscovalent bond!

42
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.

43
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).

44
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!
45
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

46
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.

47
Example of Ionic Bonding
48
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)

49
Examples of Covalent Bonding
50
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.

51
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

52
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.

53
Ionic Compounds
54
Molecular Compounds
55
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.
56
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!

57
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

58
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!

59
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!

60
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 formulathe ratio of C to H in the
  molecule.
• All ionic compounds have empirical formulas.

61
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

62
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

63
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.

64
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.

65
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.

66
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?

67
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.

68
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!

69
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

70
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

71
Determining Polarity ofMolecules
--------------------------------------------------
---------------------------
72
Hydrogen 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.
73
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

74
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

75
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.
76
Practice PercentComposition Problem

• What is the percent by mass of each element in
  Li2SO4?

77
Mole Problems

• Grams ltgt Moles
• Molecular Formula
• Stoichiometry

78
Grams ltgt 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

79
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

80
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

81
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!

82
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

83
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

84
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

85
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?

86
Concentration

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

87
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

88
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!)

89
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

90
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

91
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?

92
How many Sig Figs?

• Start counting sig figs at the first non-zero.
• All digits except place-holding zeroes are sig
  figs.

Measurement of Sig Figs
234 cm 3
67000 cm 2
_ 45000 cm 4
560. cm 3
560.00 cm 5
Measurement of Sig Figs
0.115 cm 3
0.00034 cm 2
0.00304 cm 3
0.0560 cm 3
0.00070700 cm 5
93
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

94
Rounding with addition and subtraction

• Answers are rounded to the least precise place.

95
Rounding with multiplicationand division

• Answers are rounded to the fewest number of
  significant figures.

96
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
97
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

98
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

99
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

100
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

101
Kinetics and Thermodynamics

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

102
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.

103
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)
104
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

105
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

106
Endothermic PE Diagram
If a catalyst is added?
107
Endothermic with Catalyst
The red line represents the catalyzed reaction.
108
Exothermic PE Diagram
What does it look like with a catalyst?
109
Exothermic with a Catalyst
The red line represents the catalyzed reaction.
Lower A.E. and faster reaction time!
110
Equilibrium
When the rate of the forward reaction equals the
rate of the reverse reaction.
111
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)

112
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

113
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

114
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.

115
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.

116
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.
117
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.
118
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
119
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
120
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

121
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

122
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.

123
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

124
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 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

125
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

126
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.

127
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.
128
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

129
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) ?

130
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.

131
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.
132
Alternate 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.

133
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

134
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

135
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.

136
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.

137
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!
138
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

139
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

140
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.

141
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
142
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

143
Making Voltaic Cells
More Info!!!
Create Your Own Cell!!!!
144
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.

145
You Start At The Anode
146
Make Your Own Cell!!!
147
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

148
Decomposing 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)
149
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.
150
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.
151
Organic Chemistry

• 1) Hydrocarbons
• 2) Substituted Hydrocarbons
• 3) Organic Families
• 4) Organic Reactions

152
Hydrocarbons

• Molecules made of Hydrogen and Carbon
• Carbon forms four bonds, hydrogen forms one bond
• Hydrocarbons come in three different homologous
  series
• Alkanes (single bond between Cs, saturated)
• Alkenes (1 double bond between 2 Cs,
  unsaturated)
• Alkynes (1 triple bond between 2 Cs,
  unsaturated)
• These are called aliphatic, or open-chain,
  hydrocarbons.
• Count the number of carbons and add the
  appropriate suffix!

153
Alkanes

• CH4 methane
• C2H6 ethane
• C3H8 propane
• C4H10 butane
• C5H12 pentane
• To find the number of hydrogens, double the
  number of carbons and add 2.

154
Methane
Meth- one carbon -ane alkane The simplest
organic molecule, also known as natural gas!
155
Ethane
Eth- two carbons -ane alkane
156
Propane
Prop- three carbons -ane alkane
Also known as cylinder gas, usually stored
under pressure and used for gas grills and
stoves. Its also very handy as a fuel for
Bunsen burners!
157
Butane
But- four carbons -ane alkane
Liquefies with moderate pressure, useful for gas
lighters. You have probably lit your gas grill
with a grill lighter fueled with butane!
158
Pentane
Pent- five carbons -ane alkane
Your Turn!!! Draw Hexane Draw Heptane
159
Alkenes

• C2H4 Ethene
• C3H6 Propene
• C4H8 Butene
• C5H10 Pentene
• To find the number of hydrogens, double the
  number of carbons.

160
Ethene
Two carbons, double bonded. Notice how each
carbon has four bonds? Two to the other carbon
and two to hydrogen atoms.
Also called ethylene, is used for the
production of polyethylene, which is an
extensively used plastic. Look for the PE,
HDPE (2 recycling) or LDPE (4 recycling) on
your plastic bags and containers!
161
Propene
Three carbons, two of them double bonded. Notice
how each carbon has four bonds?
If you flipped this molecule so that the double
bond was on the right side of the molecule
instead of the left, it would still be the same
molecule. This is true of all alkenes. Used to
make polypropylene (PP, recycling 5), used for
dishwasher safe containers and indoor/outdoor
carpeting!
162
Butene
This is 1-butene, because the double bond is
between the 1st and 2nd carbon from the end. The
number 1 represents the lowest numbered carbon
the double bond is touching.
This is 2-butene. The double bond is between the
2nd and 3rd carbon from the end. Always count
from the end the double bond is closest to.
ISOMERS Molecules that share the same molecular
formula, but have different structural formulas.
163
Pentene
This is 1-pentene. The double bond is on the
first carbon from the end. This is 2-pentene.
The double bond is on the second carbon from the
end. This is not another isomer of pentene.
This is also 2-pentene, just that the double bond
is closer to the right end.
164
Alkynes

• C2H2 Ethyne
• C3H4 Propyne
• C4H6 Butyne
• C5H8 Pentyne
• To find the number of hydrogens, double the
  number of carbons and subtract 2.

165
Ethyne
Now, try to draw propyne! Any isomers? Lets
see!
Also known as acetylene, used by miners by
dripping water on CaC2 to light up mining
helmets. The carbide lamps were attached to
miners helmets by a clip and had a large
reflective mirror that magnified the acetylene
flame. Used for welding and cutting applications,
as ethyne burns at temperatures over 3000oC!
166
Propyne
This is propyne! Nope! No isomers.
OK, now draw butyne. If there are any isomers,
draw them too.
167
Butyne
Well, heres 1-butyne! And heres 2-butyne!
Is there a 3-butyne? Nope! That would be
1-butyne. With four carbons, the double bond can
only be between the 1st and 2nd carbon, or
between the 2nd and 3rd carbons. Now, try pentyne!
168
Pentyne
1-pentyne 2-pentyne
Now, draw all of the possible isomers for hexyne!
169
Substituted Hydrocarbons

• Hydrocarbon chains can have three kinds of
  dingly-danglies attached to the chain. If the
  dingly-dangly is made of anything other than
  hydrogen and carbon, the molecule ceases to be a
  hydrocarbon and becomes another type of organic
  molecule.
• Alkyl groups
• Halide groups
• Other functional groups
• To name a hydrocarbon with an attached group,
  determine which carbon (use lowest possible
  number value) the group is attached to. Use di-
  for 2 groups, tri- for three.

170
Alkyl Groups
171
Halide Groups
172
Organic Families

• Each family has a functional group to identify
  it.
• Alcohol (R-OH, hydroxyl group)
• Organic Acid (R-COOH, primary carboxyl group)
• Aldehyde (R-CHO, primary carbonyl group)
• Ketone (R1-CO-R2, secondary carbonyl group)
• Ether (R1-O-R2)
• Ester (R1-COO-R2, carboxyl group in the middle)
• Amine (R-NH2, amine group)
• Amide (R-CONH2, amide group)
• These molecules are alkanes with functional
  groups attached. The name is based on the alkane
  name.

173
Alcohol
On to DI and TRIHYDROXY ALCOHOLS
Sl