AP Chemistry Review Packet Please take the time over the summer to review material learned in Chemistry 1 Honors and complete this packet. There will be a test given within the first 2 weeks on this content matter. I strongly suggest that you take the time to memorize the polyatomic ions in the packet provided Also, over the summer, purchase 2 bound composition notebooks-- college-ruled, 2 blue pens or 2 black pens and a three ring binder. Optional but not necessary: purchase a scientific calculator for your personal use during this course and a pack of 3" x 5" white index cards (pack of 100). These can be used to compose flash cards for study. I am very much looking forward to our time together next year. You will be provided with a hard copy and a digital copy of AP Chemistry. In addition, our school has purchased a digital enhancement package for each student to use. This includes but is not limited to: concept, calculation and technique videos, online practice questions with a re-teach feature as well as other helpful materials that I will share with you... in September.

Mrs. Carhart

1

Measurements

A. SI Measurements Define each. Put in correct SI units

1. Mass – quantity of matter kilogram

2. Volume –

3. Time –

4. Temperature –

5. Length –

B. Derived Units Solve. Show work.

1. Combination of fundamental units. i.e. m/s (meters per second) A sign in town gives the speed limit at 50km/hr. What is the

speed in cm/s? (answer: 1,388.9 cm/s)

2. DensityDensity = _Mass_

Volume

A piece of beeswax with a volume of 8.50 cm3 is found to have a mass of 8.06 g. What is the density of the beeswax?

What is the volume of a piece of concrete that has a mass of 8.76 g and a density of 2.85g/cm3

Cobalt is a hard metal that resembles iron in appearance. It has a density of 8.90 g/cm3. What mass would 2.00 cm3

(ml) of cobalt have?

2

C. Unit Conversions and Scientific Notation

Fill in the blanks:1. 1 L = ________ml = 1000 cm3

2. 1 m = ________nm

3. 1 m = __________ mm

4. 1 cg = ________ g

5. 100°C = _______K K = °C + 273

6. 454 g = _________kg

Write in scientific notation (Correct format 5,280 m = 5.28 x 103m)

a. 0.00025 is the diameter of a moth’s eye ___________________

b. 208,000,000,000,000,000,000,000 is the distance to another galaxy

________________

D. Uncertainty is Measurement Notes1. Rules for Significant Figures

a. Digits other than 0 ALWAYS significant .96 – 2 sig. digits

b. One or more final 0 after decimal point are significant 4.7200 – 5 sig. digits.

c. 0’s between 2 other significant digits are ALWAYS significant. 5.029 = 4 sig. digits

d. 0’s used for spacing are NOT significant 7000 – 1 sig. digit (exception: when written with decimal shown or in

scientific notation, 7.000E3 – 4 sig fig)

0.00783 – 3 sig. digit

Underline the significant digits in each of the following. How many significant digits does each value have?

35g ______ 0.004m3 ______ 3.57m _____

3

24.068kPa_____ 3.507km _____ 268k_____

0.035kg_____ 20.04080g _____ 0.246cm_____

Rules for Rounding in Calculations Noteso Addition and Subtraction--Round to the least number of

decimal placeso Multiplication and Division--Round using the least number

of significant digits

Perform the following operations and report the answer in the correct number of sig. figs.

9.44 – 2.11 = ___________ 8.30 x 2.22 = ___________

8.3 x 2.22 = ____________ 9.44 x 2.11 = ____________

E. Accuracy and Precision 1. Accuracy – how close a measurement is to the actual value

Measurement Accuracy24.3 0.1 cm18.73 0.01 cm

2. Precision – little variation with measurements made3. Error--accepted value - experimental value

4. Per cent error = |actual value - experimental value | 100% actual value

5. Qualitative vs quantitative statements

F. Dimensional Analysis Write the correct units to answer for the following problems.

1. Lab | lump | bang | bam | bog 2 | mess = | bog | lump | bog | bang | bam ___________

2. If the month of June has 30 days in it, how many seconds are in month of September? _________________

3. Mrs Carhart's turtle Bradley MUST travel 3.7 miles to get to his former home. His walk amounts to 2 inches per STEP. How many steps MUST

he take?

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4. How many seconds are there in 4 years at BRHS? Assume 180 full length days

that are 7 hours long. Set up in dimensional analysis.

Matter and Properties A. Kinds of matter 1. Homogeneous (uniform distribution) vs. Heterogeneous (more than 1 phase) 2. Solutions; solute-dissolved material, solvent- dissolving

medium

B. Physical properties1. Extensive – amount of matter present (i.e. mass, length,

volume)2. Intensive – nature of material (i.e. malleability, ductility,

conductivity, melting point, boiling point, density, specific heat)

C. Chemical Properties1. Behavior of a substance in the presence of other substances

(i.e. oxidation state, radioactivity, combustion, electronegativity, ionization energy, reactivity)

List 4 indications that a chemical change has taken place.1. ___________________________________2.___________________________________3.___________________________________4.___________________________________

How can pure substances can only separated?

How can mixtures can be separated?

Structure of the Atom Parts of the Atom Nucleus - central part of an atom; contains the protons +neutrons

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a. Very small in comparison to the entire atom. b. Held together by the strong nuclear force.

1. Atomic Number = number of protons (Z)2. Atomic Mass = number of protons + number of electrons (A)3. Number of electrons in a neutral atom = 0 (-e + +p = 0)4. If there are less electrons than protons then ion charge is positive+5. If there are fewer protons than electrons then ion charge is negative-

proton

Ionic charge

symbol neutron electron Mass number

Atomic number

12 2+ Mg 12 10 24 12

0 238

3+ 10

2- 16

32

B. Subatomic Particles Complete the table below Protons Electrons Neutrons

SymbolLocationMassCharge

C. Atomic particle research. Match the following scientists with their discovery

____ 1. J.J. Thompson

____ 2. Robert Milliken

____ 3. James Chadwick

A. discovered nucleus of an atom in gold foil experiment ;discovered protons

B. obtained 1st accurate measurement of an electron’s charge

C. discovered neutronD.

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____ 4. Earnest Rutherford D. discovered the electron

D. Isotopes Use the following table to calculate the atomic mass of titanium.

E. Unstable Nuclei and Radioactive Decay Define each. 1. Radiation—

2. Alpha Radiation—

3. Alpha particle—

4. Beta Radiation—

5. Beta Particle-- 1- charge, travels at the speed of light

7

isotope Atomic mass (amu) Relative abundance %Ti-46 45.593 8.00

___________Ti-47 46.592 7.30

___________Ti-48 47.948 73.80

___________Ti-49 48.948 5.50

____________Ti-50 49.945 5.40

____________Total Answer Here

6 Gamma Rays-

Electrons in the Atom

A. The electromagnetic Spectrum. 1. Name the major sections of the electromagnetic spectrum a.________________ b._________________

c.________________ d._________________

e.________________ f.__________________

g.________________ h.__________________

2. The above is a diagram of low frequency. In the space below, draw high frequency wavelengths. Describe the relationship between wavelength and frequency.

3. How is the flame test used to identify an element?

B. Quantum Mechanical Model

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sublevel Number of orbitals

Number of e- shape

spherical 6 dumb bell

d 10 4 leaf clover7 complex

C. Other Models of the Atom Describe the following models in detail.

a. Bohr-Rutherford (planetary) Model

b. Quantum Mechanical Model of atom. Use discoveries of Planck, de Broglie, Schrödinger, and Einstein to describe this model

D. Quantum numbers 1. (n) Principle quantum number--energy level 2. (l) Shape of the cloud – sublevels are equal to the value of the

principle quantum number 3. (m) defines each orbital direction in space of the cloud. 4. (s) describes the spin of the electron (clockwise or counter

clockwise)

9

E. Electron configuration – State each man's accomplishment (pages 156-157)

1. Auf Bau –

2. Hund –

3. Pauli Exclusion Principle –

Complete the following tables

Energy Level Number of orbitals

Maximum number of

electrons (in that energy level)

n = 1

n = 2

n = 3

n = 4

Element Configuration Orbital Filling Diagram

Electron Dot

C 1s22s22p2 _ _ __ 2s 2p

C

Ca

Al

V

Fe

U

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Periodic Table and Periodic Law A. Organization of the periodic table 1. Write the oxidation states on top of the periodic table

B.

Origin of the periodic table. Put divisions on the periodic table 1. Alkali metals #1 example

2. Transition ____

3. Lanthanides ____ 4. Alkaline earth metals _____

5. Chalcogens ______

6. Halogens ______

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7. Noble gases _____ 8. Actinides _______ 9. Nitrogen Group _____ 10. Carbon Group ______

C. Define or describe each. 1. Periodic Law--

2. Group--

3. Period--

4. Moseley achieved:

5. Mendeleyev achieved:

Ionic Compounds and Metals A. Symbols

Write the symbols for the following:Gas _____Solid _____Reversible Reaction _____Heat required _____Aqueous _____Catalyst _____Yields _____Liquid _____

B. Binary ionic compounds -- (1 metal + 1 nonmetal, end in “ide” )Example: NaCl – Sodium chloride

C. Transition metals -- (can have more than 1 charge) Some have Latin names and use the endings "ous" and "ic" to designate lesser and greater charges.

PbO2-Plumbic Oxide and PbO-Plumbous Oxide (use Roman numerals + “ide” ending) Example: V2O5 –

Vanadium (V) OxideD. Binary molecular compounds -- (made of 2 nonmetals; use

prefixes + “ide” ending)Examples: CO – carbon monoxide C2O4 – dicarbon tetroxide

Write the prefix that represents each number

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1 = mono 6 = ____ 11 = ____ 16 =____2 = di 7 = ____ 12 = ____ 17 =____ 3 = ____ 8 = ____ 13 = ____ 18 = ____4 = ____ 9 = ____ 14 = ____ 19 = ____5 = ____ 10 = ____ 15 = ____ 20 = ____

E. Polyatomic compounds

Examples: ZnCO3 – Zinc Carbonate NH4SO4 – Ammonium SulfateF. Acids

1. Ends in ide – hydro-stem-ic acidHCl = Hydrochloric Acid

2. Ends in ite – stem-ous acidH2SO3 = Sulfurous Acid

3. Ends in ate–stem ic H2SO4 = Sulfuric Acid

Write formulas for the following compounds

1. Mercury (II) Thiocynate__________________________

2. Barium Perchlorate_____________________________

3. Sodium Hypochlorite____________________________

4. Lead (II) Hypobromite___________________________

5. Tetraarsenic octatelluride________________________

6. Silver Cyanide_________________________________

7. Ammonium Carbonate__________________________

8. Copper (II) Phosphate___________________________

9. Nanosulfur Monoxide__________________________

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10. Carbon Tetafluoride__________________________

11. Potassium Dichromate__________________________

12. Lithium Thiosulfate_____________________________

13. Tin (IV) Cyanate_______________________________

14. Hydrofluoric Acid____________________________

15. Ammonium Tartrate___________________________

16. Hypochlorus Acid_____________________________

17. Aluminum Silicate_____________________________

18. Calcium Permanganate__________________________

19. Chromium (II) Oxalate__________________________

20. Potassium Thiocyanate__________________________

21. Ammonium Cyanate ____________________________

22. Sodium Thiosulfate _____________________________

23. Potassium Perchlorate __________________________

24. Nitric Acid____________________________________

25. Nitride Ion ___________________________________

26. Chloride Ion __________________________________

27. Scandium Tungstate ___________________________

28. Boron Vanadate_____________________________

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29. Lead (II) Zincate ____________________________

39. Triarsenic Hexaphosphide______________________

Write the Name for each Chemical Compound

1. Co2(S2O3)3____________________________

2. Na2 B4O7_____________________________________________

3. HCl(aq) _____________________________

4. BaSO3________________________________________________

5. Ca(NO3)2____________________________________________

6. Al2 (Cr2O7)3__________________________________________

7. Si7Se4 ______________________________

8. Sr (BrO3)2____________________________________________

9. P11N12 ______________________________

10. Mg(C2H3O2)2________________________________________

11. H2Cr2O7 (aq)__________________________

12. Li3(AsO4)______________________________

13. (NH4)3PO4____________________________________________

14. N3- _______________________________

15. Ba(NO3)2____________________________________________

16. HCH3OH(aq) _________________________

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17. La(OH)3_____________________________________________

18. Sn3(AsO3)4__________________________________________

19. Pb(CN)4_____________________________________________

20. HClO(aq)_______________________________

21. (NH4)2SO3___________________________________________

22. HClO4(aq) ______________________________

23. NH41+ _________________________________

24. Al(NO2)3 _____________________________________________

25. NH4SiO4 ______________________________

26. (NH4)2B4O7 ________________________________

27. Al(SCN)3 ________________________________

28. U(S2O3 )3 _________________________________

Write the formulas for these acids.1. Chromous Acid ________________

2. Hydrofluoric Acid ______________

3. Chlorous Acid _________________

4. Perchloric Acid ________________

5. Hydroselenic Acid ______________

6. Acetic Acid ___________________

7. Carbonic Acid _________________

8. Boric Acid ___________________

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9. Hydrosulfuric Acid _____________

10. Hypochlorous Acid _____________

11. Nitric Acid ___________________

13. Hydrotelluric Acid _____________

14. Sulfuric Acid _________________

Chemical Reactions

A. Chemical Reactions1. Reactants Products2. Balancing – Mass and Atoms are conserved. What is the Law of

Conservation of Mass and Energy ?

a. Write correct formulasb. Balance using coefficients

3. Classifying by Typea. Single replacement A + BC → AC + B (use activity series)b. Double replacement AB + CD → AD + CB (use solubility chart)c. Synthesis A + B → AB d. Decomposition AB → A + Be. Combustion CH4 + O2 → CO2 + H2O

hydrocarbon + oxygen yields carbon dioxide + water

Balance 1. AgNO3 + Na2CrO4 → Ag2CrO4 + NaNO3

2. BaCl2 + Na3PO4 → Ba3(PO4)2 + NaClC. Write complete chemical equations and balance. Identify type of reaction.

17

1. Solid elemental Iron combines with solid Sulfur to produce Iron (II) Sulfide.

2. Gaseous Sulfur Dioxide combines with Water (l) to produce Sulfurous Acid

3. Heat is applied to Lead (IV) Oxide to produce elemental Lead + Oxygen gas.

4. Mrs. Carhart dropped a small piece of Sodium into a beaker of water to create

a violent reaction. It produced Hydrogen gas and Sodium Hydroxide.

D. Write the single replacement reaction + balance. Then write complete + net ionic equations. Remember: A must be higher than B on the

activities series for A to replace B. See sample below.

Balanced: Cu(s) + 2Ag (NO3)(aq) Cu (NO3)2(aq) + 2Ag(s)Complete Ionic Equation:Cu0(s) + 2NO31-(aq) + 2Ag1+(aq) Cu2+ (aq) +2NO31-(aq) + 2Ag0(s)Net Ionic Equation:Cu0(s) + 2Ag1+ (aq) Cu2+(aq) + 2Ag0(s)

Write correct equations

1. Balanced: Fe(s) + Pb(NO3)2(aq) Fe(NO3)2 + Pb

2. Complete Ionic:

3. Net Ionic:

E. Double replacement reactions. Write balanced equation and predict the precipitate using the solubility chart. Then write net ionic equations. If both products are aqueous, no reaction is visible (NR). See example below.

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Balanced equationBaCl2 + Na2(CrO4) -> 2NaCl + Ba(Cr2O4 )(s) Complete ionicBa2+

(aq)+ 2Cl1-(aq) + 2Na1+

(aq)+ Cr2O42-

(aq) -> BaCr2O4 (s) + 2Na1-(aq)

2Cl1-(aq)

Net ionicBa2+

(aq)+ Cr2O42-

(aq) -> BaCr2O4 (s)

1. Balanced: NaCl (aq) + AgNO3(aq) → NaNO3 + AgCl(s)

2. Complete Ionic:

3. Net Ionic:

F. Combustion Reactions. Balance in the following manner if a hydrocarbon is used.CxHy + O2 → xCO2 + y/2H2O Formula can be used to solve for C and H. Then balance the oxygenWrite a balanced equation for the complete combustion of:

1. Glycerol C3H8O3 + 5O2 → 3CO2 + 4H2O

2. Octane C8H18

3. Hexane C8H14

Chemical Quantities-The Mole A. Define each of the following:

a. Mole-

b. Avogadro’s number-

c. molar volume-

d. molecular formula-

19

e. empirical formula-

f. per cent composition-

g. hydrate-

B. Describe the relationship between Avogadro’s number and one mole of ANY substance.

C. State whether the formula is molecular or empirical.

1. S2Cl2_______________

2. Na2SO4_____________

3. C17H19NO3___________ 4. C9H10O4 _____________ 5. C5H10O5_____________

6. H2SO4 ______________

D. Calculate the percent composition of Capsaicin (Hot Peppers)?

1. C18H27NO3--formula for what makes hot peppers hot!

2. Which of the following compounds has the highest iron content? Determine per cent composition of the iron of each.

a. FeCl2

b. Fe(C2H3O2)3

20

c. Fe(OH)2

E. What is the empirical formula of the compound that is 28.2% K, 25.6% Cl and 46.2% O

F. Molecular Formula 1. What is the molecular formula of the compound that has a molar mass of 160 grams and an empirical formula of SO3?

2. What is the molecular formula of a compound 94.1% O and 5.9% H, molar mass = 34g

G. Find the molar mass of each of the following compounds and name the particle type (atom, molecule, formula unit or ion)

1. Li2S ______________________________

2. Ca(OH)2 ___________________________

3. Na2CrO4 ___________________________

4. U ________________________________

5. C6H12O6 ___________________________

H. Solve for moles, liters, grams or particles.

1. How many grams are in 9.45 mol of dinitrogen trioxide? (718.2 g)

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2. What is the volume at STP of 9.1 mol of dinitrogen trioxide gas?

3. Convert 25.0 grams of Copper (Cu) into atoms.

4. Convert 8.80x1023 atoms of Platinum (Pt) into grams of Pt.

5. Assuming STP, how many moles are in these volumes?

A. 67.2L SO2 (3 mol)

B. 4.00 « 103L C2H6

5. What is a hydrate? Name the following hydrates: (see page 351)

A. MgSO4 « 7H2O_____________________________

B. Na2B4O7 « 10H2O___________________________

C. Ba(OH)2 « 8H2O___________________________

Stoichiometry A. Define each of the following:

a. Limiting Reagent-

b. Excess Reactant-

c. Percent yield-

d. Theoretical yield-

B. Problems to Solve:

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How many grams of copper (I) sulfide could be produced from 9.90 g of copper (I) chloride reacting with an excess of hydrogen sulfide gas?(write a balanced equation first)

How many grams of H2 can be produced from the reaction of 72.0 g of sodium with an excess of water?

If 5.00 g of copper metal react with a solution containing 20.0 g of AgNO3 to produce silver metal, which reactant is limiting?

Cu(s) + 2AgNO3(aq) Cu(NO3)2(aq) + 2Ag(s)

Calcium carbonate can be decomposed by heating. Mrs. Carhart wants to know what is the per cent yield of this reaction if 24.8g of CaCO3 is heated to give 13.1g CaO? Actual yield = 13.1grams

CaCO3(s) CaO(s) + CO2(g)

Ionic Bonding and Metals A. Electronegativity

What is the charge when EACH of the following elements loses its valence electrons?

a. aluminumb. lithium c. bariumd. strontium

Write electron configurations for the following and comment on the result.

a. N3-

b. O2-

c. F-

d. Ne

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Which of the following pairs of elements WILL NOT form ionic compounds?

a. Sulfur and Oxygenb. Sodium and Calciumc. Sodium and Sulfurd. Oxygen and Chlorine

Calculate electronegativity difference for the following: (subtract)

a. Ba – O (3.44 – 0.89 ) = 2.55

b. Ca – Cl

c. Br – Rb

d. Li – Br

e. F – F

f. H – Cl

B. Bond types (ionic, covalent, metallic) NOTES1. Ionic Bond

A. Electronegativity greater than 1.67B. Crystalline solid at room temperatureC. High melting points/ boiling pointsD. Conducts electricity in molten or liquid states

2. Covalent BondA. Electronegativity less than 1.67 polar or nonpolar (often

nonpolar bonds with an electronegativity of 0 to 0.4)B. Exists as gases or liquids at room temperatureC. Low melting/boiling pointsD. Has weak intermolecular forces called: van der Waals, dipole-

dipole interactions, dispersion forces, hydrogen bonding3. Metallic Bond

A. + Metal ions surrounded by a sea of valence electronsB. Valence electrons are mobile and travel/conducts electricityC. Crystalline solids—body-centered cubic, face-centered cubic,

hexagonal packCovalent Bonding

Distinguish between: polar covalent and nonpolar covalent bonds:

Lewis dot structure

24

Electron Configurations

Multiple Choice1. The electron configuration of a fluoride ion F1-

a. Is22s22p5 b. same as Ne c. 1s22s22p63s1 d. same as K

2. Which of these elements does NOT exist as a diatomic molecule?

a. Neb. Fc. Hd. I

3. Which one of the following compounds is NOT covalenta. BrClb. LiClc. HCld. S2Cl2

4. A diatomic molecule with a triple covalent bond is:a. N2 b. F2 c. H2

5. To gain a noble gas configuration a sulfur atom must:a. gain two electronsb. lose 1 electronc. lose 2 electronsd. Mrs. Carhart says it gains 3 electrons

6. An ionic compound is:a. a saltb. held together by ionic bondsc. composed of anions and cationsd. all of the above

7. A metallic bond is a bond between:

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a. valence electrons and + charged metal ionsb. ions of two different metalsc. metal and a non metald. Mrs Carhart says "none of the above"

8. Metals are good conductors of electricity because they:a. are hardb. are ductilec. contain mobile valence electronsd. bend easily

Define each of the following:

13. Single –1 shared pair of electrons

14. Double – ____________________

15. Triple –_______________________

16. Pi Bond-________________________

17. Coordinate Covalent Bond __________________________Identify the bond type below

Complete the following table:Molecul

eLewis dot Structure

Shape Structural Formula

Polar/NonpolarCompound

1. CH4 tetrahedral

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Molecule

Lewis dot Structure

Shape Structural Formula

Polar/NonpolarCompound

2. CO2 linear

2. BeF2

linear

3. NH3 trigonal

4. C6H6

ring

5. C2H2

XIII. Periodic Trends Chapter 6A. Important relationships in the periodic table. For above answers, choose the term INCREASE (↑), DECREASE (↓)REMAINS THE SAME (S)

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Property Periods(moving across period from

left to right)

Groups(moving vertically down a

group)

Atomic weightAtomic numberAtomic radius

Positive Oxidation state

Negative oxidation state

Ionic radius (for positive oxidation

number)Electronegativity

Ionization energy

Answer the following questions for atomic number 92 using your Periodic Table

1. Mass number

2. Number of protons

3. Number of electrons

4. Number of neutrons

5. Electron configuration

6. Electron dot

7. Orbital filling diagram

8. More or less metallic than Neodymium

9. Ionization energy greater or less than Nd

10. Density greater or less than Nd

28

11. Metal, nonmetal, metalloid, noble gas

12. Solid, liquid or gas at room temperature

13. What period is it in?

14. name of element

15. charges of the U ion

16. Noble Gas that has the same electron configuration as the Nd ion

17. Atomic radius is bigger or smaller than Nd

18. More or less electronegative than Nd

19. Ionization energy is higher or lower than for Nd

20. Is U likely to lose electrons?

21. U ion is larger or smaller than Nd ion?

22. Is U ion larger or smaller than U atom?

1. Acetate C2H3O2-1

2. Amide NH2-1

3. Arsenate AsO4-3

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4. Arsenite AsO3-3

5. Azide N3-1

6. Benzoate C6H5COO-1

7. Bisulfite HSO3-1

8. Borate BO3-3

9. Bromate BrO3-1

10. Carbonate CO3-2

11. Chlorate ClO3-1

12. Chlorite ClO2-1

13. Chromate CrO4-2

14. Chromite CrO2-1

15. Citrate C6H5O7-3

16. Cyanate OCN-1

17. Cyanide CN-1

18. Dichromate Cr2O7-2

19. Dithionate S2O4-2

20. Dihydrogen phosphate H2PO4-1

21. Formate HCOO-1

22. Glutamate C5H8NO4-1

23. Hexachloroplatinate PtCl6-2

24. Hexacyanoferrate Fe(CN)6-3

25. Hydrogen carbonate HCO3-1

26. Hydroxide OH-1

27. Hypobromite BrO-1

28. Hypochlorite ClO-1

29. Hypoiodite IO-1

30. Hydrogen oxalate HC2O4-1

31. Hydrogen sulfate HSO4-1

32. Hydrogen sulfite HSO3-1

33. Imide NH-1

34. Iodate IO3-1

35. Iodite IO2-1

36. Lactate C3H5O3-1

37. Manganate MnO4-2

38. Molybdate MoO4-2

39. Nitrate NO3-1

40. Nitrite NO2-1

41. Orthosilicate SiO4-4

42. Oxalate C2O4-2

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43. Perbromate BrO4-1

44. Perchlorate ClO4-1

45. Periodate IO4-1

46. Permanganate MnO4-1

47. Peroxide O2-2

48. Phosphate PO4-3

49. Phosphite PO3-3

50. Monohydrogen phosphateHPO4-2

51. Selenate SeO4-2

52. Selenite SeO32-

53. Selenocyanate SeCN-1 54. Silicate SiO3

-2

55. Sorbate C6H7O2-1

56. Sulfate SO4-2

57. Sulfite SO3-2

58. Tartrate C4H4O6-2

59. Tellurate TeO4-2

60. Tetraborate B4O7-2

61. Thiocyanate SCN-1

62. Thiosulfate S2O3-2

63. Triodide I3-1

64. Tungstate WO4-2

65. Vanadate VO3-1

66. Zincate ZnO2-2

Ammonium NH4 1+ Hydronium H3O 1+ Sulfate SO4 2-

31

Activity Series of MetalsReactivity decreases down a group

Name symbol

1 Lithium Li2 Potassiu

mK

3 Barium Ba4 Calcium Ca5 Sodium Na6 Magnesiu

mMg

7 Aluminum

Al

8 Zinc Zn9 Iron Fe10

Nickel Ni

11

Tin Sn

12

Lead Pb

13

Hydrogen H*

14

Copper Cu

15

Mercury Hg

16

Silver Ag

17

Gold Au

32

Some Useful Conversion Factors

TEMPERATURE: SI = Kelvin (K)F = ( 1.8ºC ) + 32 = Fahrenheit (°F)K = C + 273.15C = ( F-32 ) x 5/9 = Celsius (°C) VOLUME SI = Meter3 (m3)1 meter3 = 106centimeters3 = 103liter1 cm3 = 1 millimeter = 10-

3liter1 gallon = 4 quarts = 8 pints = 32 fl.oz. 1 Liter = 1.057 quarts = 1000cm3

ENERGY SI = Joule (J)1 calorie = 4.184 Joules1 Joule = 107ergs1000 calories = 1 kilocalorie (kcal)

MASS SI = Kilogram (kg)1 kilogram = 2.205 pound1 pound = 16 ounces = 453.6 grams1 ton = 2000 pounds

LENGTH SI = Meter (m)1 mile = 5280 feet = 1.609 kilometers1 meter = 39.37 inches1 angstrom (Å) = 10-10meter

PRESSURE SI = Pascal (Pa)1 atmosphere =

COLORS ABSORBED IN THEVISIBLE REGION

Wavelength Color Absorbed 400nm

ultraviolet400-450 violet450-490 blue490-550 green550-580 yellow580-650 orange650-700 red 700nm infrared

PHYSICAL CONSTANTS:

Avogadro’s number N = 6.0221 1023

Gas constant R = 0.082058 L atm/mol K

Gravitational acceleration g = 9.8066 m/s

Planck’s constant h = 6.6261 10-34 J s

Speed of light in vacuum c = 2.9979 108 m/s

Specific Heat of Water Cp= 4.184 J/g°C

Density of Water = 1.000 g/cm3 @ 4°C

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