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Final exam review

  • To insert your company logo on this slide

  • From the Insert Menu

  • Select “Picture”

  • Locate your logo file

  • Click OK

  • To resize the logo

  • Click anywhere inside the logo. The boxes that appear outside the logo are known as “resize handles.”

  • Use these to resize the object.

  • If you hold down the shift key before using the resize handles, you will maintain the proportions of the object you wish to resize.


What to bring to the exam
What to Bring to the Exam

  • Two #2 pencils with good erasers.

  • Your brain. Please don’t leave it at home. 


    • PDA’s and cell. Cell phones must be turned OFF and be left somewhere other than where you are sitting.

    • Headphones may not be on your person during the exam.

    • Calculators will not be allowed (they won’t be needed).

    • Food or any drinks.

(c) 2006, Mark Rosengarten

How to prepare
How To Prepare

  • DO NOT CRAM. Get your studying done with by the night before. Get a good night’s sleep and have breakfast the morning of the exam.

  • Use notes and information on the blog.

  • Actively participate in any and all review classes and activities offered by your teacher.

  • Study vocabulary. Identify key words and use flash cards to help you remember what the meaning of those words are and the concepts behind them.

(c) 2006, Mark Rosengarten

Outline for review
Outline for Review

1) The Atom (Electron Config)

2) Matter (Phases, Types, Changes)

3) Bonding (Periodic Table, Ionic, Covalent)

4) Compounds (Formulas, Reactions, IMAF’s)

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)

10) ) The Atom (Nuclear)

(c) 2006, Mark Rosengarten

The atom
The Atom

1) Nucleons

2) Isotopes

3) Electron Configuration

4) Development of the Atomic Model

(c) 2006, Mark Rosengarten


  • 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.

(c) 2006, Mark Rosengarten


  • 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.

(c) 2006, Mark Rosengarten

Weight average atomic mass
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)

(c) 2006, Mark Rosengarten

Most common isotope
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 doesn’t always work, but it usually does. Use it for the Regents exam.

(c) 2006, Mark Rosengarten

Electron configuration
Electron Configuration

  • Basic Configuration

  • Valence Electrons

  • Electron-Dot (Lewis Dot) Diagrams

  • Excited vs. Ground State

  • What is Light?

(c) 2006, Mark Rosengarten

Basic configuration
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)

(c) 2006, Mark Rosengarten

Valence electrons
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.

(c) 2006, Mark Rosengarten

Electron dot diagrams
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.


(c) 2006, Mark Rosengarten

Example dot diagrams
Example Dot Diagrams

Carbon can also have this dot diagram, which it

has when it forms organic compounds.

(c) 2006, Mark Rosengarten

Excited vs ground state
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 doesn’t 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

(c) 2006, Mark Rosengarten

What is light
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.


(c) 2006, Mark Rosengarten

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.

(c) 2006, Mark Rosengarten

Development of the atomic model
Development of the Atomic Model

  • Thompson Model

  • Rutherford Gold Foil Experiment and Model

  • Bohr Model

  • Quantum-Mechanical Model

(c) 2006, Mark Rosengarten

Thompson model
Thompson Model

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

(c) 2006, Mark Rosengarten

Rutherford model
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


(c) 2006, Mark Rosengarten

Bohr model
Bohr Model

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

(c) 2006, Mark Rosengarten

Quantum mechanical model
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 can’t 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.

(c) 2006, Mark Rosengarten


1) Properties ofPhases

2) Types of Matter

3) Phase Changes

(c) 2006, Mark Rosengarten

Properties of phases
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)

(c) 2006, Mark Rosengarten


The positive and

negative ions

alternate in the

ionic crystal lattice

of NaCl.

(c) 2006, Mark Rosengarten


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


(c) 2006, Mark Rosengarten


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.

(c) 2006, Mark Rosengarten

Types of matter
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.

(c) 2006, Mark Rosengarten


  • 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.

(c) 2006, Mark Rosengarten


  • 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!

(c) 2006, Mark Rosengarten


  • 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.

(c) 2006, Mark Rosengarten

Phase changes
Phase Changes

  • Phase Change Types

  • Phase Change Diagrams

  • Heat of Phase Change

  • Evaporation

(c) 2006, Mark Rosengarten

Phase change types
Phase Change Types

(c) 2006, Mark Rosengarten

Phase change diagrams
Phase Change Diagrams

AB: Solid Phase

BC: Melting (S + L)

CD: Liquid Phase

DE: Boiling (L + G)

EF: Gas Phase

Notice how temperature remains constant during a phase change? That’s because the PE is changing, not the KE.

(c) 2006, Mark Rosengarten

Heat of phase change
Heat of Phase Change

  • How many joules would it take to melt 100. g of H2O (s) at 0oC?

  • q=mHf = (100. g)(334 J/g) = 33400 J

  • How many joules would it take to boil 100. g of H2O (l) at 100oC?

  • q=mHv = (100.g)(2260 J/g) = 226000 J

(c) 2006, Mark Rosengarten


  • 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 liquid’s 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).

(c) 2006, Mark Rosengarten


1) The Periodic Table

2) Ions

3) Ionic Bonding

4) Covalent Bonding

5) Metallic Bonding

(c) 2006, Mark Rosengarten

The periodic table
The Periodic Table

  • Metals

  • Nonmetals

  • Metalloids

  • Chemistry of Groups

  • Electronegativity

  • Ionization Energy

(c) 2006, Mark Rosengarten


  • 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)

(c) 2006, Mark Rosengarten


  • 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)

(c) 2006, Mark Rosengarten


  • 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).

(c) 2006, Mark Rosengarten

Chemistry of groups
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

(c) 2006, Mark Rosengarten

Group 1 alkali metals
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

(c) 2006, Mark Rosengarten

Group 2 alkaline earth metals
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

(c) 2006, Mark Rosengarten

Groups 3 11 transition metals
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

  • Cu+1 = copper (I) Cu+2 = copper (II)

  • Compounds containing these metals can be colored.

(c) 2006, Mark Rosengarten

Group 17 halogens
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.

(c) 2006, Mark Rosengarten

Group 18 noble gases
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.

(c) 2006, Mark Rosengarten

Diatomic molecules
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: 2H2O  2 H2 + O2

(c) 2006, Mark Rosengarten


  • An atom’s attraction to electrons in a chemical bond.

  • F has the highest, at 4.0

  • Fr has the lowest, at 0.7

  • If two atoms that are different in EN (END) from each other by 1.7 or more collide and bond (like a metal atom and a nonmetal atom), the one with the higher electronegativity will pull the valence electrons away from the atom with the lower electronegativity to form a (-) ion. The atom that was stripped of its valence electrons forms a (+) ion.

  • If the two atoms have an END of less than 1.7, they will share their unpaired valence electrons…covalent bond!

(c) 2006, Mark Rosengarten

Ionization energy
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.

(c) 2006, Mark Rosengarten

Final exam review

  • 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).

(c) 2006, Mark Rosengarten

Metal ions cations
Metal Ions (Cations)

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!

  • Na: 2-8-1

  • Na+1: 2-8

  • Ca: 2-8-8-2

  • Ca+2: 2-8-8

  • Al: 2-8-3

  • Al+3: 2-8

(c) 2006, Mark Rosengarten

Nonmetal ions anions
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

(c) 2006, Mark Rosengarten

Ionic bonding
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.

(c) 2006, Mark Rosengarten

Example of ionic bonding
Example of Ionic Bonding

(c) 2006, Mark Rosengarten

Covalent bonding
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”)

(c) 2006, Mark Rosengarten

Examples of covalent bonding
Examples of Covalent Bonding

(c) 2006, Mark Rosengarten

Metallic bonding
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.

(c) 2006, Mark Rosengarten


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

(c) 2006, Mark Rosengarten

Types of compounds
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 [H+A- (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., don’t conduct.

(c) 2006, Mark Rosengarten

Ionic compounds
Ionic Compounds

(c) 2006, Mark Rosengarten

Molecular compounds
Molecular Compounds

(c) 2006, Mark Rosengarten

Network solids
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.

(c) 2006, Mark Rosengarten

Formula writing
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.

  • Na+1 and Cl-1 = NaCl

  • Ca+2 and Br-1 = CaBr2

  • Al+3 and O-2 = Al2O3

  • Zn+2 and PO4-3 = Zn3(PO4)2

  • Try these problems!

(c) 2006, Mark Rosengarten

Formulas to write
Formulas to Write

  • Ba+2 and N-3

  • NH4+1 and SO4-2

  • Li+1 and S-2

  • Cu+2 and NO3-1

  • Al+3 and CO3-2

  • Fe+3 and Cl-1

  • Pb+4 and O-2

  • Pb+2 and O-2

(c) 2006, Mark Rosengarten

Formula naming
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!

(c) 2006, Mark Rosengarten

The stock system
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!

(c) 2006, Mark Rosengarten

Empirical formulas
Empirical Formulas

  • Ionic formulas: represent the simplest whole number mole ratio of elements in a compound.

  • Ca3N2 means a 3:2 ratio of Ca ions to N ions in the compound.

  • Many molecular formulas can be simplified to empirical formulas

    • Ethane (C2H6) can be simplified to CH3. This is the empirical formula…the ratio of C to H in the molecule.

  • All ionic compounds have empirical formulas.

(c) 2006, Mark Rosengarten

Molecular formulas
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

(c) 2006, Mark Rosengarten

Types of chemical reactions
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

(c) 2006, Mark Rosengarten


  • Two elements combine to form a compound

  • 2 Na + O2 Na2O

  • Same reaction, with charges added in:

    • 2 Na0 + O20 Na2+1O-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.

(c) 2006, Mark Rosengarten


  • A compound breaks down into its original elements.

  • Na2O  2 Na + O2

  • Same reaction, with charges added in:

    • Na2+1O-2 2 Na0 + O20

  • O-2 is oxidized (loses electrons), is the reducing agent

  • Na+1 is reduced (gains electrons), is the oxidizing agent

  • Electrons are transferred from the O-2 to the Na+1.

  • No spectator ions, there are only two elements here.

(c) 2006, Mark Rosengarten

Single replacement
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 K+1Cl-1 Ca+2Cl2-1 + 2 K0

  • Ca0 is oxidized (loses electrons), is the reducing agent

  • K+1 is reduced (gains electrons), is the oxidizing agent

  • Electrons are transferred from the Ca0 to the K+1.

  • Cl-1 is the spectator ion, since it’s charge doesn’t change.

(c) 2006, Mark Rosengarten

Double replacement
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 Ag+1 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 Na+1 and NO3-1 ions remain dissolved throughout the reaction, they are the spectator ions.

  • How do identify the precipitate?

(c) 2006, Mark Rosengarten

Identifying the precipitate
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.

(c) 2006, Mark Rosengarten

Balancing chemical reactions
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!

(c) 2006, Mark Rosengarten

Reactions to balance
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

(c) 2006, Mark Rosengarten

Attractive forces
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

(c) 2006, Mark Rosengarten

Determining polarity of molecules
Determining Polarity ofMolecules


(c) 2006, Mark Rosengarten

Hydrogen bond attractions
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.

(c) 2006, Mark Rosengarten

Math of chemistry
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

(c) 2006, Mark Rosengarten

Formula mass
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

(c) 2006, Mark Rosengarten

Percent composition
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. Don’t forget to take atomic mass to the nearest tenth! This is a problem for you to try.

(c) 2006, Mark Rosengarten

Practice percent composition problem
Practice PercentComposition Problem

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

(c) 2006, Mark Rosengarten

Mole problems
Mole Problems

  • Grams <=> Moles

  • Molecular Formula

  • Stoichiometry

(c) 2006, Mark Rosengarten

Grams moles
Grams <=> Moles

  • How many grams will 3.00 moles of NaOH (40.0 g/mol) weigh?

  • 3.00 moles X 40.0 g/mol = 120. g

  • How many moles of NaOH (40.0 g/mol) are represented by 10.0 grams?

  • (10.0 g) / (40.0 g/mol) = 0.250 mol

(c) 2006, Mark Rosengarten

Molecular formula
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

(c) 2006, Mark Rosengarten


  • 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

(c) 2006, Mark Rosengarten

Gas laws
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!

(c) 2006, Mark Rosengarten

Gas law problem 1
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

(c) 2006, Mark Rosengarten

Gas law problem 2
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

(c) 2006, Mark Rosengarten

Gas law problem 3
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

(c) 2006, Mark Rosengarten


  • 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?

(c) 2006, Mark Rosengarten


  • Molarity

  • Parts per Million

  • Percent by Mass

  • Percent by Volume

(c) 2006, Mark Rosengarten


  • 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

(c) 2006, Mark Rosengarten

Parts per million
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 (it’s OK!)

(c) 2006, Mark Rosengarten

Percent by mass
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%

(c) 2006, Mark Rosengarten

Percent by volume
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%

(c) 2006, Mark Rosengarten

Sig figs and rounding
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?

(c) 2006, Mark Rosengarten

How many sig figs
How many Sig Figs?

  • Start counting sig figs at the first non-zero.

  • All digits except place-holding zeroes are sig figs.

(c) 2006, Mark Rosengarten

What precision
What Precision?

  • A number’s 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

(c) 2006, Mark Rosengarten

Rounding with addition and subtraction
Rounding with addition and subtraction

  • Answers are rounded to the least precise place.

(c) 2006, Mark Rosengarten

Rounding with multiplication and division
Rounding with multiplicationand division

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

(c) 2006, Mark Rosengarten

Metric conversions
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

(c) 2006, Mark Rosengarten


  • 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

(c) 2006, Mark Rosengarten

Solving for q
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

(c) 2006, Mark Rosengarten

Solving for m
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

(c) 2006, Mark Rosengarten

Solving for d t
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

(c) 2006, Mark Rosengarten

Kinetics and thermodynamics
Kinetics and Thermodynamics

1) Reaction Rate

2) Heat of Reaction

3) Potential Energy Diagrams

4) Equilibrium

5) Le Châtelier’s Principle

6) Solubility Curves

(c) 2006, Mark Rosengarten

Reaction rate
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 catalystmakes a reaction go faster by removing steps from the mechanism and lowering the activation energy without getting used up in the process.

(c) 2006, Mark Rosengarten

Heat of reaction
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)

(c) 2006, Mark Rosengarten

Potential energy diagrams
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

(c) 2006, Mark Rosengarten

Making a pe diagram
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



(c) 2006, Mark Rosengarten

Endothermic pe diagram
Endothermic PE Diagram

If a catalyst is added?

(c) 2006, Mark Rosengarten

Endothermic with catalyst
Endothermic with Catalyst

The red line represents the catalyzed reaction.

(c) 2006, Mark Rosengarten

Exothermic pe diagram
Exothermic PE Diagram

(c) 2006, Mark Rosengarten

What does it look like with a catalyst?

Exothermic with a catalyst
Exothermic with a Catalyst

The red line represents the catalyzed reaction. Lower A.E. and faster reaction time!

(c) 2006, Mark Rosengarten


When the rate of the forward reaction equals the rate of the reverse reaction.

(c) 2006, Mark Rosengarten

Examples of equilibrium
Examples of Equilibrium

  • Solution Equilibrium: when a solution is saturated, the rate of dissolving equals the rate of precipitating.

    • NaCl (s)  Na+1 (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)

(c) 2006, Mark Rosengarten

Le ch telier s principle
Le Châtelier’s 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

(c) 2006, Mark Rosengarten

Steps to relieving stress
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

(c) 2006, Mark Rosengarten


  • 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.

(c) 2006, Mark Rosengarten

Solubility curves
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.

(c) 2006, Mark Rosengarten


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.

(c) 2006, Mark Rosengarten


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.

(c) 2006, Mark Rosengarten


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

(c) 2006, Mark Rosengarten

How ionic solutes dissolve in water
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

(c) 2006, Mark Rosengarten

Acids and bases
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

(c) 2006, Mark Rosengarten

Formulas naming and properties of acids
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

(c) 2006, Mark Rosengarten

Properties of acids
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.

(c) 2006, Mark Rosengarten

Naming of acids
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

(c) 2006, Mark Rosengarten

Formula writing of acids
Formula Writing of Acids

  • Acids formulas get written like any other. Write the H+1 first, then figure out what the negative ion is based on the name. Cancel out the charges to write the formula. Don’t forget the (aq) after it…it’s only an acid if it’s in water!

  • Hydrosulfuric acid: H+1 and S-2 = H2S (aq)

  • Carbonic acid: H+1 and CO3-2 = H2CO3 (aq)

  • Chlorous acid: H+1 and ClO2-1 = HClO2 (aq)

  • Hydrobromic acid: H+1 and Br-1 = HBr (aq)

  • Hydronitric acid:

  • Hypochlorous acid:

  • Perchloric acid:

(c) 2006, Mark Rosengarten

Formulas naming and properties of bases
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)  Na+1 (aq) + OH-1 (aq)

  • Properties of Bases

  • Naming of Bases

  • Formula Writing of Bases

(c) 2006, Mark Rosengarten

Properties of bases
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”.

(c) 2006, Mark Rosengarten

Naming of bases
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.

(c) 2006, Mark Rosengarten

Formula writing of bases
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 = Ca+2 and OH-1 = Ca(OH)2 (aq)

  • Potassium hydroxide = K+1 and OH-1 = KOH (aq)

  • Lead (II) hydroxide = Pb+2 and OH-1 = Pb(OH)2 (aq)

  • Lead (IV) hydroxide = Pb+4 and OH-1 = Pb(OH)4 (aq)

  • Lithium hydroxide =

  • Copper (II) hydroxide =

  • Magnesium hydroxide =

(c) 2006, Mark Rosengarten


  • H+1 + 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) 

(c) 2006, Mark Rosengarten

Final exam review

  • 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.

(c) 2006, Mark Rosengarten


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.

(c) 2006, Mark Rosengarten

Alternate theories
Alternate Theories

  • Arrhenius Theory: acids and bases must be in aqueous solution.

  • Alternate Theory: Not necessarily so!

    • Acid: proton (H+1) donor…gives up H+1 in a reaction.

    • Base: proton (H+1) acceptor…gains H+1 in a reaction.

  • HNO3 + H2O H3O+1 + NO3-1

    • Since HNO3 lost an H+1 during the reaction, it is an acid.

    • Since H2O gained the H+1 that HNO3 lost, it is a base.

(c) 2006, Mark Rosengarten

Oxidation and reduction
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

(c) 2006, Mark Rosengarten

Oxidation numbers
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

(c) 2006, Mark Rosengarten

Identifying ox rd si species
Identifying OX, RD, SI Species

  • Ca0 + 2 H+1Cl-1Ca+2Cl-12 + H20

  • Oxidation = loss of electrons. The species becomes more positive in charge. For example, Ca0Ca+2, so Ca0 is the species that is oxidized.

  • Reduction = gain of electrons. The species becomes more negative in charge. For example, H+1H0, so the H+1 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.

(c) 2006, Mark Rosengarten


  • Ca0 + 2 H+1Cl-1Ca+2Cl-12 + H20

  • Since Ca0 is being oxidized and H+1 is being reduced, the electrons must be going from the Ca0 to the H+1.

  • Since Ca0 would not lose electrons (be oxidized) if H+1 weren’t there to gain them, H+1 is the cause, or agent, of Ca0’s oxidation. H+1 is the oxidizing agent.

  • Since H+1 would not gain electrons (be reduced) if Ca0 weren’t there to lose them, Ca0 is the cause, or agent, of H+1’s reduction. Ca0 is the reducing agent.

(c) 2006, Mark Rosengarten

Writing half reactions
Writing Half-Reactions

  • Ca0 + 2 H+1Cl-1Ca+2Cl-12 + H20

  • Oxidation: Ca0Ca+2 + 2e-

  • Reduction: 2H+1 + 2e-H20

The two electrons lost by Ca0 are gained by the two H+1 (each H+1 picks up an electron).


(c) 2006, Mark Rosengarten

Practice half reactions
Practice Half-Reactions

  • Don’t 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:

(c) 2006, Mark Rosengarten

Balancing half reactions
Balancing Half-Reactions

  • Ca0 + Fe+3 Ca+2 + Fe0

    • Ca’s charge changes by 2, so double the Fe.

    • Fe’s charge changes by 3, so triple the Ca.

    • 3 Ca0 + 2 Fe+3 3 Ca+2 + 2 Fe0

  • Try these:

  • __Na0 + __H+1 __Na+1 + __H20

    • (hint: balance the H and H2 first!)

  • __Al0 + __Cu+2 __Al+3 + __Cu0

(c) 2006, Mark Rosengarten

Activity series
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.

(c) 2006, Mark Rosengarten

Metal activity
Metal Activity

3 K0 + Fe+3Cl-13


  • 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 K+1Cl-1


(c) 2006, Mark Rosengarten

Voltaic cells
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

(c) 2006, Mark Rosengarten

Making voltaic cells
Making Voltaic Cells







(c) 2006, Mark Rosengarten

How it works
How It Works

  • The Zn0 anode loses 2 e-, which go up the wire and through the load. The Zn0 electrode gets smaller as the Zn0 becomes Zn+2 and dissolves into solution. The e- go into the Cu0, where they sit on the outside surface of the Cu0 cathode and wait for Cu+2 from the solution to come over so that the e- can jump on to the Cu+2 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.

Since Zn is listed above Cu, Zn0 will be oxidized when it reacts with Cu+2. The reaction: Zn + CuSO4  ZnSO4 + Cu

(c) 2006, Mark Rosengarten

You start at the anode
You Start At The Anode

(c) 2006, Mark Rosengarten

Make your own cell
Make Your Own Cell!!!

(c) 2006, Mark Rosengarten

Electrolytic cells
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

(c) 2006, Mark Rosengarten

Decomposing alkali metal compounds
Decomposing AlkaliMetal Compounds

2 NaCl  2 Na + Cl2

The Na+1 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)

(c) 2006, Mark Rosengarten

Decomposing water
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.

(c) 2006, Mark Rosengarten


The Ag0 is oxidized to Ag+1 when the (+) end of the battery strips its electrons off.

The Ag+1 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.

(c) 2006, Mark Rosengarten

Organic chemistry
Organic Chemistry


2) Substituted Hydrocarbons

3) Organic Families

4) Organic Reactions

(c) 2006, Mark Rosengarten


  • 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 C’s, saturated)

    • Alkenes (1 double bond between 2 C’s, unsaturated)

    • Alkynes (1 triple bond between 2 C’s, unsaturated)

  • These are called aliphatic, or open-chain, hydrocarbons.

  • Count the number of carbons and add the appropriate suffix!

(c) 2006, Mark Rosengarten


  • CH4 = methane

  • C2H6 = ethane

  • C3H8 = propane

  • C4H10 = butane

  • C5H12 = pentane

  • To find the number of hydrogens, double the number of carbons and add 2.

(c) 2006, Mark Rosengarten


Meth-: one carbon

-ane: alkane

The simplest organic molecule, also known as natural gas!

(c) 2006, Mark Rosengarten


Eth-: two carbons

-ane: alkane

(c) 2006, Mark Rosengarten


Prop-: three carbons

-ane: alkane

Also known as “cylinder gas”, usually stored under pressure and used for gas grills and stoves. It’s also very handy as a fuel for Bunsen burners!

(c) 2006, Mark Rosengarten


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!

(c) 2006, Mark Rosengarten


Pent-: five carbons

-ane: alkane

Your Turn!!!

Draw Hexane:

Draw Heptane:

(c) 2006, Mark Rosengarten


  • C2H4 = Ethene

  • C3H6 = Propene

  • C4H8 = Butene

  • C5H10 = Pentene

  • To find the number of hydrogens, double the number of carbons.

(c) 2006, Mark Rosengarten


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!

(c) 2006, Mark Rosengarten


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!

(c) 2006, Mark Rosengarten


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.

(c) 2006, Mark Rosengarten


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.

(c) 2006, Mark Rosengarten


  • C2H2 = Ethyne

  • C3H4 = Propyne

  • C4H6 = Butyne

  • C5H8 = Pentyne

  • To find the number of hydrogens, double the number of carbons and subtract 2.

(c) 2006, Mark Rosengarten


Now, try to draw propyne! Any isomers? Let’s see!

Also known as “acetylene”, used by miners by dripping water on CaC2 to light up mining helmets. The “carbide lamps” were attached to miner’s 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!

(c) 2006, Mark Rosengarten


This is propyne! Nope! No isomers.

OK, now draw butyne. If there are any isomers, draw them too.

(c) 2006, Mark Rosengarten


Well, here’s 1-butyne!

And here’s 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!

(c) 2006, Mark Rosengarten




Now, draw all of the possible isomers for hexyne!

(c) 2006, Mark Rosengarten

Substituted hydrocarbons
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.

(c) 2006, Mark Rosengarten

Alkyl groups
Alkyl Groups

(c) 2006, Mark Rosengarten

Halide groups
Halide Groups

(c) 2006, Mark Rosengarten

Organic families
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.

(c) 2006, Mark Rosengarten



(c) 2006, Mark Rosengarten

Di and tri hydroxy alcohols
Di and Tri-hydroxy Alcohols

(c) 2006, Mark Rosengarten

Positioning of functional group
Positioning of Functional Group

PRIMARY (1o): the functional group is bonded to a carbon that is on the end of the chain.

SECONDARY (2o): The functional group is bonded to a carbon in the middle of the chain.

TERTIARY (3o): The functional group is bonded to a carbon that is itself directly bonded to three other carbons.

(c) 2006, Mark Rosengarten

Organic acid
Organic Acid

These are weak acids. The H on the right side is the one that ionized in water to form H3O+. The -COOH (carboxyl) functional group is always on a PRIMARY carbon.

Can be formed from the oxidation of primary alcohols using a KMnO4 catalyst.

(c) 2006, Mark Rosengarten


Aldehydes have the CO (carbonyl) groups ALWAYS on a PRIMARY carbon. This is the only structural difference between aldehydes and ketones.

Formed by the oxidation of primary alcohols with a catalyst. Propanal is formed from the oxidation of 1-propanol using pyridinium chlorochromate (PCC) catalyst.*

(c) 2006, Mark Rosengarten


Ketones have the CO (carbonyl) groups ALWAYS on a SECONDARY carbon. This is the only structural difference between ketones and aldehydes.

Can be formed from the dehydration of secondary alcohols with a catalyst. Propanone is formed from the oxidation of 2-propanol using KMnO4 or PCC catalyst.*

(c) 2006, Mark Rosengarten


Ethers are made of two alkyl groups surrounding one oxygen atom. The ether is named for the alkyl groups on “ether” side of the oxygen. If a three-carbon alkyl group and a four-carbon alkyl group are on either side, the name would be propyl butyl ether. Made with an etherfication reaction.

(c) 2006, Mark Rosengarten


Esters are named for the alcohol and organic acid that reacted by esterification to form the ester. If the alcohol was 1-propanol and the acid was hexanoic acid, the name of the ester would be propyl hexanoate. Esters contain a COO (carboxyl) group in the middle of the molecule, which differentiates them from organic acids.

(c) 2006, Mark Rosengarten


  • Component of amino acids, and therefore proteins, RNA and DNA…life itself!

  • - Essentially ammonia (NH3) with the hydrogens replaced by one or more hydrocarbon chains, hence the name “amine”!

(c) 2006, Mark Rosengarten


Synthetic Polyamides: nylon, kevlar

Natural Polyamide: silk!

For more information on polymers, go here.

(c) 2006, Mark Rosengarten

Organic reactions
Organic Reactions

  • Combustion

  • Fermentation

  • Substitution

  • Addition

  • Dehydration Synthesis

    • Etherification

    • Esterification

  • Saponification

  • Polymerization

(c) 2006, Mark Rosengarten


  • Happens when an organic molecule reacts with oxygen gas to form carbon dioxide and water vapor. Also known as “burning”.

(c) 2006, Mark Rosengarten


  • Process of making ethanol by having yeast digest simple sugars anaerobically. CO2 is a byproduct of this reaction.

  • The ethanol produced is toxic and it kills the yeast when the percent by volume of ethanol gets to 14%.

(c) 2006, Mark Rosengarten


  • Alkane + Halogen  Alkyl Halide + Hydrogen Halide

  • The halogen atoms substitute for any of the hydrogen atoms in the alkane. This happens one atom at a time. The halide generally replaces an H on the end of the molecule.

    C2H6 + Cl2 C2H5Cl + HCl

    The second Cl can then substitute for another H:

    C2H5Cl + HCl  C2H4Cl2 + H2

(c) 2006, Mark Rosengarten


  • Alkene + Halogen  Alkyl Halide

  • The double bond is broken, and the halogen adds at either side of where the double bond was. One isomer possible.

(c) 2006, Mark Rosengarten


  • Alcohol + Alcohol  Ether + Water

  • A dehydrating agent (H2SO4) removes H from one alcohol’s OH and removes the OH from the other. The two molecules join where there H and OH were removed.

Note: dimethyl ether and diethyl ether are also produced from this reaction, but can be separated out.

(c) 2006, Mark Rosengarten


  • Organic Acid + Alcohol  Ester + Water

  • A dehydrating agent (H2SO4) removes H from the organic acid and removes the OH from the alcohol. The two molecules join where there H and OH were removed.

(c) 2006, Mark Rosengarten


The process of making soap from glycerol esters (fats).

Glycerol ester + 3 NaOH  soap + glycerol

Glyceryl stearate + 3 NaOH  sodium stearate + glycerol

The sodium stearate is the soap! It emulsifies grease…surrounds globules with its nonpolar ends, creating micelles with - charge that water can then wash away. Hard water replaces Na+ with Ca+2 and/or other low solubility ions, which forms a precipitate called “soap scum”.

Water softeners remove these hardening ions from your tap water, allowing the soap to dissolve normally.

(c) 2006, Mark Rosengarten


  • A polymer is a very long-chain molecule made up of many monomers (unit molecules) joined together.

  • The polymer is named for the monomer that made it.

    • Polystyrene is made of styrene monomer

    • Polybutadiene is made of butadiene monomer

  • Addition Polymers

  • Condensation Polymers

  • Rubber

(c) 2006, Mark Rosengarten

Addition polymers
Addition Polymers

Joining monomers together by breaking double bonds

Polyvinyl chloride (PVC): vinyl siding, PVC pipes, etc.

Vinyl chloride polyvinyl chloride

n C2H3Cl  -(-C2H3Cl-)-n

Polytetrafluoroethene (PTFE, teflon):


n C2F4 -(-C2F4-)-n

(c) 2006, Mark Rosengarten

Condensation polymers
Condensation Polymers

Condensation polymerization is just dehydration synthesis, except instead of making one molecule of ether or ester, you make a monster molecule of polyether or polyester.

(c) 2006, Mark Rosengarten


The process of toughing rubber by cross-linking the polymer strands with sulfur is called...

(c) 2006, Mark Rosengarten


(c) 2006, Mark Rosengarten

Natural radioactivity
Natural Radioactivity

  • Alpha Decay

  • Beta Decay

  • Positron Decay

  • Gamma Decay

  • Charges of Decay Particles

  • Natural decay starts with a parent nuclide that ejects a decay particle to form a daughter nuclide which is more stable than the parent nuclide was.

(c) 2006, Mark Rosengarten

Alpha decay
Alpha Decay

  • The nucleus ejects two protons and two neutrons. The atomic mass decreases by 4, the atomic number decreases by 2.

  • 23892U 

(c) 2006, Mark Rosengarten

Beta decay
Beta Decay

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

  • 146C 

(c) 2006, Mark Rosengarten

Positron decay
Positron Decay

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

  • 5326Fe 

(c) 2006, Mark Rosengarten

Gamma decay
Gamma Decay

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

(c) 2006, Mark Rosengarten

Charges of decay particles
Charges of Decay Particles

(c) 2006, Mark Rosengarten

Half life

  • Half life is the time it takes for half of the nuclei in a radioactive sample to undergo decay.

  • Problem Types:

    • Going forwards in time

    • Going backwards in time

    • Radioactive Dating

(c) 2006, Mark Rosengarten

Going forwards in time
Going Forwards in Time

  • How many grams of a 10.0 gram sample of I-131 (half-life of 8 days) will remain in 24 days?

  • #HL = t/T = 24/8 = 3

  • Cut 10.0g in half 3 times: 5.00, 2.50, 1.25g

(c) 2006, Mark Rosengarten

Going backwards in time
Going Backwards in Time

  • How many grams of a 10.0 gram sample of I-131 (half-life of 8 days) would there have been 24 days ago?

  • #HL = t/T = 24/8 = 3

  • Double 10.0g 3 times: 20.0, 40.0, 80.0 g

(c) 2006, Mark Rosengarten

Radioactive dating
Radioactive Dating

  • A sample of an ancient scroll contains 50% of the original steady-state concentration of C-14. How old is the scroll?

  • 50% = 1 HL

  • 1 HL X 5730 y/HL = 5730y

(c) 2006, Mark Rosengarten

Nuclear power
Nuclear Power

  • Artificial Transmutation

  • Particle Accelerators

  • Nuclear Fission

  • Nuclear Fusion

(c) 2006, Mark Rosengarten

Artificial transmutation
Artificial Transmutation

  • 4020Ca + _____ -----> 4019K + 11H

  • 9642Mo + 21H -----> 10n + _____

  • Nuclide + Bullet --> New Element + Fragment(s)

  • The masses and atomic numbers must add up to be the same on both sides of the arrow.

(c) 2006, Mark Rosengarten

Particle accelerators
Particle Accelerators

  • Devices that use electromagnetic fields to accelerate particle “bullets” towards target nuclei to make artificial transmutation possible!

  • Most of the elements from 93 on up (the “transuranium” elements) were created using particle accelerators.

  • Particles with no charge cannot be accelerated by the charged fields.

(c) 2006, Mark Rosengarten

Nuclear fission
Nuclear Fission

  • 23592U + 10n 9236Kr + 14156Ba + 3 10n + energy

  • The three neutrons given off can be reabsorbed by other U-235 nuclei to continue fission as a chain reaction

  • A tiny bit of mass is lost (mass defect) and converted into a huge amount of energy.

(c) 2006, Mark Rosengarten

Chain reaction
Chain Reaction

(c) 2006, Mark Rosengarten

Nuclear fusion
Nuclear Fusion

  • 21H + 21H 42He + energy

  • Two small, positively-charged nuclei smash together at high temperatures and pressures to form one larger nucleus.

  • A small bit of mass is destroyed and converted into a huge amount of energy, more than even fission.

(c) 2006, Mark Rosengarten

The end

(c) 2006, Mark Rosengarten