Chem

EDGE ECE Specialist

CHEMISTRY

EDGE ECE REVIEW SPECIALIST 1

Chemistry

CCHHEEMMIISSTTRRYY

Chemistry is a branch of science which deals with the study of matter and the changes it undergoes. A. Branches of chemistry

Organic Chemistry Inorganic Chemistry Physical Chemistry Analytical Chemistry Biochemistry

I. MATTER Matter is anything that occupies space and has mass. It is composed of tiny particles called atoms. There are presently 106 different kinds of atoms (elements) in which each of these are represented by a symbol.

Law of Conservation of Mass Mass can neither be created nor be destroyed. Law of Conservation of Energy Energy can not be created nor destroyed. It can only be

transformed from one form to another. Law of Definite Composition A pure compound is always made up of same constituent

elements combined in a definite proportion by weight. Law of Multiple Proportions When two elements react to form more than one compound,

the different weights of one that combine with a fixed weight of the other are in the ratio of small whole numbers.

SCHEMATIC DIAGRAM-CLASSIFICATION OF MATTER

A. Physical States Solid has definite size and shape Liquid has definite volume but takes the shape of the container Gas has neither definite shape nor definite volume

B. Properties Physical Properties those that can be measured without changing the basic identity of the

substance (e.g. color, density, odor, boiling point). Chemical Properties those that describe how a substance may react to form other

substances (e.g. flammability, tendency to rust).


Chemistry

Intrinsic Properties are properties of the substance that are independent of the shape and size of the substance. (e.g. temperature, pressure, etc.).

Extrinsic Properties are properties of the substance that are related to its size and shape. (e.g. volume, mass, weight, etc.)

C. Changes Which Matter Undergoes Physical Change involves changing one or more physical properties of a sample of matter

without changing its composition (e.g. evaporation, cutting of wire, crystallization, tearing of paper).

Chemical Change results in the change in composition of matter (e.g. burning of paper, rusting of iron).

D. Composition

Mixtures Pure Substance

Heterogeneous Mixtures marble, concrete, wood

Compounds salt, water, carbon dioxide gas

Homogeneous Mixtures sugar solution, pure air, metal alloys

Elements hydrogen gas, gold, mercury, neon gas

II. CHEMICAL FORMULAS

A. The Atom Basic building block of the universe; has the following major components:

Subatomic particles Charges Mass Electric Charge Unit mass in kg

proton electron neutron

positive negative neutral

1.0073 amu 0.00055 amu 1.0087 amu

1.602x10-19 C -1.602x10-19 C

1.673x10-27 kg 9.11x10-31 kg 1.675x10-27 kg

Atomic number equal to the number of protons of an element. Atomic mass equal to the combined masses of protons and neutrons. Neutral atom number of protons is equal to the number of electrons. Positively charged atom if there are more protons than electrons. Negatively charged atom if there are more electrons than protons.

Example: What is the number of neutron in one atom of 5626

Fe ? Answer: The number of protons is 26. The number of protons and neutrons (mass number) is 56. Thus, the number of neutrons is 56 26, or 30.

B. Atomic Weight Equal to the average of the isotopic masses weighted according to the naturally occurring

abundance of the isotopes of the element. Expressed relative to the value of exactly 12 amu for a carbon-12 atom.


Chemistry

Example: What is the atomic weight of argon (Ar) given the following percentage of abundance in nature: 99.60% 40Ar, mass is 39.962 amu 0.337% 36Ar, mass is 35.968 amu 0.063% 38Ar, mass is 37.963 amu. Answer: Atomic weight of argon: = (0.996x39.962) + (0.00337x35.968) + (0.00063x37.963) = 39.947 amu

C. Formula Weight Is used for compounds that are made up of ions and have primarily ionic bonding. Is

convenient as it can be used for both ionic and covalent bonding. D. Molecular Weight Is used for compounds that are composed of molecules and have primarily covalent

compound. Will be used only for covalent compounds which consists of molecules like sucrose C12H22O11, ethyl alcohol C2H5OH, and Carbon Monoxide CO.

Example: Calculate the formula weight of water, H2O. Answer: Since there are 2 atoms of hydrogen and 1 atom of oxygen in a formula unit of water, then the formula weight is H = 2 x 1 = 2 O = 1 x 16 = 16 Formula Weight = 18 amu

E. Mole Amount of a substance which contains 6.022 x 1023 particles (Avogadros number) of matter;

(the world particle can mean atom, molecule, or ion). Equal to the gram molecular mass of a substance

Example: How many moles and atoms are there in 100 g of argon? Answer: The molecular mass of argon is 39.948 g/mole. The number of moles of argon in the sample is:

= 1mol100g39.948g

= 2.50 mol Ar. The number of atoms of argon is simply

=236.022 10 atoms2.50 moles

mole

=1.51 x 1023 atoms

F. Formulas and Formula Masses A chemical formula indicates the relative number of atoms of each element in a substance. If the numbers are the smallest possible integral values that express the relative number of

atoms, the formula is called an empirical formula. The formula mass is the sum of the masses of every atom in a substance as indicated in its

formula.


Chemistry

Example: A compound consists of 30.4% nitrogen and the rest oxygen. What is its empirical formula?

Answer: mole N = 1moleN30.4g N14gN

= 2.17 mole N

mole O = (100 30.4)g O 1mol O16.0g O

= 4.35 mole O

The smallest mole ratio of nitrogen to oxygen is 1:2. Thus, the empirical formula is NO2.

G. Molecular Formulas and Molecular Masses A molecular formula is similar to the empirical formula expect that it expresses the actual

number of atoms of each element in one molecule of a substance. If a molecular formula is used, the corresponding formula mass is called a molecular mass.

Example: A compound with molecular mass equal to 60.0 has the following percent composition: C = 40.0%, H = 6.67%, O = 53.3%. What is its molecular formula? Answer: Atomic weight of argon: = (0.996x39.962) + (0.00337x35.968) + (0.00063x37.963) = 39.947 amu

II. CHEMICAL COMPOUNDS

A. Types of Chemical Compounds Ionic Compounds metal and nonmetal ions held together by electrostatic forces of

attraction. Molecular Compounds nonmetal elements held together by covalent bonding. B. Composition of Chemical Compounds It is the components of compounds and their relative proportions in a given sample.

Example: What are the masses of carbon and hydrogen in 50.0 g of methane (CH4)? Answer: The molecular mass of methane is (4)1 + (1)12, or 16. The percentage of C is 12/16 or 75.0%. The percentage of hydrogen is 4/16 or 25.0%. For a 50.0g sample the mass of hydrogen is = 0.25 (50.0g) = 12.5g And that of carbon is = 0.75 (50g) = 37.5g

C. Oxidation State It is the number of electron an atom can donate, accept, or share with other atoms to from a

compound. The common oxidation states of some elements are the following:


Chemistry

Elements H F, Cl, Br, I O, S N, P C, Si B, Al Alkali metals

Alkaline earth

metals

Oxidation state 1, -1 -1 -2 -3, +5 -4, +4 +3 +1 +2

D. Nomenclature of Inorganic Compounds Binary compounds (composed of two elements)

Rules for Metals and Nonmetals: 1. The unmodified name of the metal is written followed by the name of the nonmetal, which

ends in ide. 2. For transition metals, a suffix ous is added for the lower state while ic for the higher one. 3. If the Stock System is used, the oxidation number of the metal is written in Roman Numeral

right after the unmodified name of the metal.

Compound Name Compound Name

Fe2S3 ferrous sulfide AlF3 aluminum fluoride

BaO barium oxide Cr2O3 chromium(III) oxide

Cu2O copper(I) oxide ZnS zinc sulfide

CaF2 calcium fluoride SrO strontium oxide

Na2S sodium sulfide MgCl2 magnesium chloride

Rules for two Nonmetals: 1. Prefixes are written to indicate the relative number of atoms of an element in a compound. 2. A suffix ide is added at the end.


BCl3 boron trichloride SF6 sulfur hexaoxide

CCl4 carbon tetrachloride PCl3 phosphorus trichloride

CO carbon monoxide PCl5 phosphorus pentachloride

NO2 nitrogen dioxide B2Br4 diboron tetrabromide

N2O dinitrogen oxide SO2 sulfir dioxide

Ternary compounds composed of more than two elements, usually a polyatomic ion and an

element. Naming them is by order: positive first, negative second


NH4Cl ammonium chloride Na4PO4 sodium phosphate

KC2H3O2 potassium acetate NaCN sodium cyanide


Chemistry

Mg(NO3)2 magnesium nitrate BaC2O4 barium oxalate

NaHCC3 sodium bicarbonate KMnO4 potassium permanganate

K2CrO4 potassium chromate Na2S2O3 sodium thiosulfate

Binary acids a prefix hydro- and a suffix -ic are added to the base name of the nonmetallic

element, then the word acid.

Name Formula

Hydrofluoric acid Hydrochloric acid Hydrobromic acid Hydroiodic acid Hydrosulfuric acid

HF HCl HBr HI H2S

Ternary acids composed of hydrogen, a nonmetal, and oxygen. Naming them depends on

the of oxygen present in the acid, usually with the lesser number ending with ous and with the greater number ending in ic; others follow the name of their polyatomic ions.

Name Formula Name Formula Name Formula

Nitric acid HNO3 Hypochloric acid HClO Phosphorous acid H3PO3

Nitrous acid HNO2 Chlorous acid HClO2 Phosphoric acid H3PO4

Sulfuric acid H2SO4 Chloric acid HClO3 Boric acid H3BO3

Sulfurous acid H2SO3 Perchloric acid HClO4 Carbonic acid H2CO3

Acetic acid HC2H3O2 Oxalic acid H2C2O4 Silicic acid H2SiO3

III. CHEMICAL REACTIONS AND STOICHIOMETRY

A. Types of Chemical Reactions Combination/synthesis formation of a compound of complex substance through the

reaction of two elements of simpler substances. Examples: 2H2 + O2 2H2O + heat 2C7H6O2 + 15O2 14CO2 + 6H2O + heat CaO + CO2 CaCO3 Decomposition/analysis breakdown of a compound into other compounds and/or elements. Examples: 2H2O 2H2 + O2 2HgO 2Hg + O2 CaCO3 CaO + CO2 Single replacement/single displacement the more reactive element replaces the less

reactive element in a compound. Examples: 2Mg + TiCl4 2MgCl + Ti Zn + CuSO4 ZnSO4 + Cu


Chemistry

Double replacement/double displacement (also called metathesis) exchange of patterns to form an insoluble salt.

Examples: BaCl2 + NaSO4 2NaCl + BaSO4 Na2CO3 + Ba(OH)2 BaCO3 + 2NaOH B. Chemical Equation and Stoichiometry Chemical equation representation of a chemical reaction; reactants are written on the left

side, products at the right side of the arrow. Law of conservation of matter in a chemical reaction, total mass of reactants equals total

mass of the products. Stoichiometric coefficients numbers written before a substance in balancing an equation. C. Limiting Reactant It is the reactant that restricts or controls the amount of product that will be produced. D. Percent Yield Not all reactions proceed to 100% completion, that is not all reactants are consumed to yield

the desired product. Some reactants undergo side-reactions to produce unintended products (the by-products).

The percent yield is defined as the ratio of the actual yield over the theoretical yield times 100.

IV. GAS LAWS

A. Boyles Law For a fixed amount of gas at constant temperature, gas volume is inversely proportional to gas

pressure.

1 2

2 1

v Pv P

Example: A certain was occupying a volume of 10L at 720 mm Hg. At constant temperature, the gas was compressed resulting to a pressure of 800 mm Hg. What was the new volume of the gas?

Answer: 12 1

2

P 720 mm Hgv v 10 L 9 LP 800 mm Hg

B. Charles Law For a fixed amount of gas at constant pressure the gas volume is directly proportional to the

absolute temperature of the gas (i.e., in Kelvin scale).

1 1

2 2

v Tv T

Example: Air inside a 5 L frictionless piston at 25C was heated up to 50C. What was the new volume of the air?


Chemistry

Answer: 22 1

1

T 273 50 Cv v L 5.42 LT 273 25 C

5

C. Gay-Lussacs Law For a fixed amount of gas at constant volume, gas pressure is directly proportional to gas

temperature.

1 1

2 2

P TP T

Example: Oxygen gas at 30C and 10 atm was further pressurized to 15 atm by heating the tank. What was the new temperature of the oxygen gas?

Answer: 22 11

P 15atmT T 30 273 K 454.5K 181.5 CP

10atm

D. Avogadros Law At a fixed pressure and temperature, the volume of a gas is directly proportional to the amount

of gas.

1 1

2 2

v nv n

At STP or standard temperature and pressure (0C and 1 atm) the volume of a mole of

gas is 22.4 L. E. The Ideal Gas Law At low pressure and high temperature, all gases follow the above gas laws. The combination of

all the above laws is called the Ideal Gas Law and it follows the following equation:

PV = nRT, where R is the ideal gas constant equal to L atm0.0821mole K

The equation can also be expressed as: 1 1 2 2

1 2

PV P V... nR

T T

Example: Carbon dioxide occupies a volume of 3L at 1.5 atm and 47C. How many moles of carbon dioxide are there? If it is cooled down to 30C and subjected to a pressure of 2 atm, what is the new volume of the gas? Answer: PV = nRT


Chemistry

1 1 2 21 2

1 22 1

2 1

PV (1.5 atm)(3 L)n 0.17 moleRT Latm0.0821 (47 273)K

moleKP V P VT T

PT 1.5 atm (30 27)KV V 3 L 2.13 LP T 2 atm (47 273)K

F. Daltons Law of Partial Pressures The total pressure of a gaseous mixture is equal to the sum of the partial pressure of the

individual gases that make up the mixture. The partial pressure of a component gas is simply the pressure that gas is exerting on a

container as if it were alone.

Example: Air at standard atmospheric pressure is typically 78.084% nitrogen, 20.946% oxygen, 0.934%?argon, and 0.036% carbon dioxide. What are the partial pressures of each gas in mm HG? Answer: Conversion: 1 atm = 760 mm Hg Partial pressure of: Nitrogen: 0.78084 x 760 = 593.44 mm Hg Oxygen: 0.20946 x 760 = 159.19 mm Hg Argon: 0.00934 x 760 = 7.10 mm Hg Carbon Dioxide: 0.00036 x 760 = 0.27 mm Hg

G. Grahams Law of Effusion Effusion the escape of a gas through an orifice or hole. The rate of effusion of a gas is

inversely proportional to the square root of its molecular weight.

1 2

2 1

r MWr MW

V. THERMOCHEMISTRY

A. Terminology Heat (q) an energy transfer due to a temperature difference. Work (w) a form of energy transfer between a system and its surroundings in the form of

compression of expansion of gas. Internal energy (U) the total energy attributed to the particles of matter and their interactions

within a system; composed of thermal energy (energy associated with random molecular motion) and chemical energy (energy associated with chemical bonds and intermolecular forces).

Enthalpy (H) a thermodynamic function defined by H = U + PV. At constant temperature and pressure, the change in enthalpy, H, is simply the heat of reaction.

Heat reaction (qrxn) heat exchange in a system when theres a chemical reaction at constant temperature.


Chemistry

Heat capacity (c) the amount of heat required to raise the temperature of an object or substance by one degree; usually expressed in J/C.

Specific heat/molar heat capacity (cp) heat capacity per unit mass of a substance at constant pressure.

Latent heat of fusion (Lf) heat absorbed to melt a substance at constant temperature. Latent heat of vaporization (Lv) heat required to change a substance from its liquid phase

to its gaseous phase at constant temperature. B. Calorimetry (heat measurement) Change in temperature q = heat capacity x temperature change = CT = mass of object x specific heat x temperature change = mCpT, for water, specific heat is Cp = 1.0 cal/gC

Example: A mass of 50g of copper (specific heat = 0.093 cal/gC) at 30C is heated up to 100C. How much heat was absorbed by the copper? Answer: q = mCpT

q = cal50g 0.093 100 C 30 Cg C

q = 325.5 cal

Phase change 1. Solid liquid q = mLv, for water heat of fusion is: Lf = 80 cal/g 2. Liquid gas q = mLf, for water heat of vaporization is: Lv = 540 cal/g

Example: If 100g of ice (Cp = specific heat = 0.5 cal/gC) at -5C is converted to steam (Cp = specific heat = 0.5 cal/gC) at 110C, how much heat is required? Answer: heat involved in: temperature change (-5C to 0C) q1 = mCpT = 100g (0.5 cal/gC)(5C) = 250 cal phase change (ice to water) q2 = mLv = 100g (80 cal/g) = 8000 cal temperature change (0C to 100C) q3 = mCpT = 100g (1cal/gC)(100C) = 10,000 cal phase change (water to steam) q4 = mLf = 100g(540cal/g) = 54,000 cal Total heat = q1 + q2 + q3 + q4 = (250 + 8000 + 10000 + 54000) cal = 72250 cal

Chemical reaction 1. Endothermic process a reaction wherein heat is absorbed by the system, indicated by a

positive change of enthalpy. 2. Exothermic process a reaction wherein heat is released by the system, indicated by a

negative change of enthalpy. 3. Hess Law states that if a reaction is carried out in a series of steps, the enthalpy of

reaction, H, is equal to the sum of the enthalpy changes for the individual steps.


Chemistry

VI. THE ATOMIC STRUCTURE

The Photoelectric Effect discovered by H. Hertz in 1888; described the emission of electrons from metal surfaces when struck with light of appropriate frequency.

Photon proposed by Einstein in 1905; it means a particle of light consisting a particular amount (a quantum) of energy. When it collides with an electron, it gives up its entire energy to the electron.

Bohrs Theory of a Hydrogen Atom introduced by Niels Bohr in 1913; states that 1) an electron in an atom can only be in a certain allowed places, and 2) when it is in one of these allowed places it possesses a certain amount of energy.

Wave-Particle Duality proposed by Louis de Broglie; states that small particles of matter may at times display wave-like properties.

The Uncertainty Principle established by Werner Heisenberg; states that it is impossible to know the precise location and velocity of an electron at the same time.

The Schrdinger Wave Equation formulated by Erwin Schrdinger; describes the wave-mechanical model of electrons in an atom.

Orbital a region in an atom where the electron charge density or the probability of finding an electron is high.

Quantum Numbers the three integral numbers needed to solve the equation of wave mechanics.

1. Principal quantum number (n) refers to the average distance of the orbital from the nucleus.

n = 1, 2, 3 2. Orbital angular quantum number (l) refers to the shape of the orbital. The specific orbital

types are s (sharp), p (principal), d (diffuse), and f (fundamental). Its value depends on the principal quantum number.

l = 0, 1, 2 (n-1)

3. Magnetic quantum number (ml) refers to the spatial orientation of the orbital. Its value

depends on the angular quantum number. 4. Spin quantum number (ms) refers to the spin of the electron, sometimes regarded as the

fourth quantum number. The value can be either +1/2 or -1/2.

Rules for Assigning Electrons Orbitals 1. Paulis Exclusion Principle no two electrons in an atom can have the same set of

quantum numbers. 2. Hunds Rule pf Maximum Multiplicity whenever orbitals of equal energy are available,

electrons occupy these orbitals singly before any pairing of electrons.

The Aufbau Process a method of writing the probable electron configuration of the elements in the order of increasing atomic number.

Example: Give the electron configuration of gallium, Ga, with atomic number 31. Answer: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1


Chemistry

VII. THE PERIODIC TABLE

Periodic table graphical arrangement of the elements in order of increasing atomic numbers such that elements with similar properties are arranged in vertical columns.

Periodic Law when all the elements are arranged in order of increasing atomic numbers, elements with similar properties will occur at periodic intervals.

Family / Group a vertical columns of elements in the periodic table that provides the number of valence electrons. e.g., Family 1-A: H, Li, Na, K, Rb, Cs, Fr.

Series / Period horizontal row in the periodic table that provides the number of the last main energy level. E.g., Series 3: Na, Mg, Al, Si, P, S, Cl, Ar.

A. The Main Groups in the Periodic Table

Group Number Family Name

Group 1A: The Alkali Metals

Group 2A: The Alkaline Earth Metals

Group 3A: The Aluminum Group

Group 3A: The Boron Family

Group 4A: The Tin and Lead Family

Group 4A: The Carbon and Silicon Family

Group 5A: The Nitrogen Family

Group 6A: The Oxygen Family

Group 7A: The Halogens

Group 8A: The Noble Gases

Group B: The Transition Metals

Classifications of Elements

1. Metals good conductors of heat and electricity, ductile, malleable, and shiny. All metals are

solid at room temperature with the exception of liquid mercury. They are the elements, except Hydrogen, that are on the left side of the border line including the Lanthanide and Actinide metals.

2. Non-metals poor conductors of heat and electricity, brittle, not shiny, with more varied physical properties than metals. They are all the elements on the right side of the border line such as S, Br, and Ar.

3. Metalloids with properties that fall between those of metals and non-metals. They are the elements that are above and below the borderline plus elements of group 4A such as Al, C, and As.

Trends in the Periodic Table

1. Electronegativity ability of an atom to attract electron. 2. Ionization energy energy needed to remove an electron from an atom. 3. Electron affinity energy released when an electron is added to an atom.

VIII. CHEMICAL BONDING


Chemistry

Chemical Bonds forces of attraction that exist between atoms. Lewis Symbols and Lewis Structures consist of a chemical symbol with dots placed around it to

represent the valence electrons. Types of Chemical Bonds

1. Ionic or electrovalent bond formed by the transfer of electron from a metallic element to a non-metallic element.

e.g. (1) NaCl (2) Fe2O3 2. Covalent bond formed by the sharing of electrons between two or more non-metallic elements.

polar covalent bond unequal sharing of electrons e.g. (1) H2O (2) NH3 non-polar covalent bond equal sharing of electrons e.g. (1) CH4 (2) O2 (3) N2 dative bond/coordinate covalent bond the pair of electrons shared between atoms is

donated by only one atom. e.g. (1) NH3BF3 (2) NH4

3. Metallic bond force of attraction that exists within elemental metals (e.g., all metallic elements). 4. Double-Triple Bond If there are two pairs of electrons between two atoms, it is called a

DOUBLE BOND, and if there are three pairs it is called a TRIPLE BOND.

Resonance the use of two or more Lewis structures to represent a particular molecule. e.g. (1) SO2 (2) O3 Isomers substances that have the same molecular formulas but differ in their structures and in

their properties. e.g. Molecular Formula = C2H6O

Ethanol = C2H5OH dimethyl ether = CH3OCH3

VIII. LIQUIDS, SOLIDS, AND IMF

Comparison of Liquids and Solids

State of Matter Volume/Shape Density Compressibility

Motion of Molecules

Liquid

Has a definite volume; its shape follows the shape

of its container

High Only slightly compressible

The molecules slide past one another freely;

liquids are fluid.

Solid Has a definite shape and volume

High (generally, solids are

denser than liquids)

Incompressible Vibrate about fixed positions


Chemistry

Surface Tension the amount of energy required to increase the surface by unit area. Liquids with strong intermolecular forces of attraction exhibit greater surface tension.

Viscosity resistance to flow. Liquids which have strong intermolecular forces of attraction are less fluid than those which have weak forces of attraction. Liquid sugar is thick and flows very slowly.

Intermolecular forces attractive forces that exist between molecules. Types of IMF

1. Van der Waals Forces very weak intermolecular forces that exist between non-polar molecules.

e.g. (1) CH4 (2) H2 2. Dipole-Dipole Forces forces that act between polar molecules. e.g. (1) HCl (2) H2O 3. Hydrogen Bonding interaction between the hydrogen atoms bonded to an atom of a very

electronegative element (F, N, O). e.g. (1) H2O (2) NH3

IX. SOLUTIONS AND THEIR PHYSICAL PROPERTIES

Two components of a solution 1. Solute dissolved substance, present in lesser quantity 2. Solvent dissolved medium, present in greater quantity

Types of solutions according to the solubility of solute 1. Saturated solution a solution containing the maximum amount of solute that can be dissolved

by the solvent at a given temperature. 2. Unsaturated solution a solution containing less solute than the solvent can dissolve at a given

temperature. 3. Supersaturated solution a solution containing more solute than the solvent can dissolve.

Factors affecting solubility of solute

1. Nature of solute and solvent 2. Temperature 3. Size of particles, Surface area 4. Pressure (solids and liquids are not affected)

Henrys Law: the solubility of a gas in a liquid is proportional to the pressure of the gas over the solution.

Expressing Solution Concentration

1. Percent solution a. % by mass = (mass of solute / mass of solution) x 100%

Example: A sample of 0.892 g of naphthalene (C10H8) is dissolved in 54.6 g of benzene (C6H6). What is the percent by mass of naphthalene in this solution?

Answer: percent by mass 0.892g 100%

0.892g 54.6g1.61%


Chemistry

b. % by volume = (volume of solute/volume of solution) x 100% c. ppm = (mass of volume of solute/mass or volume of solution) x 1,000,000 d. proof = twice the % of alcohol in solution 2. Mole Fraction (X) the no. of moles of a component divided by the total number of moles of all

components in the solution.

AA

A B

BB

A B

nX

n nn

Xn n

Note: A = solute B = solvent

Example: Determine the mole fractions of both substances in a solution containing 26.0 of NaCl and 125.0 g of water. Answer:

A

2

A

2B

2

26.0gNaCl58.5g/ moleNaCl

X26.0gNaCl 125.0gwater

58.5 / moleNaCl 18g/ moleH O

X 0.06

125.0gwater18g /moleH O

X26.0gNaCl 125.0gwater

58.5 /moleNaCl 18g /moleH O

BX 0.94

3. Molarity (M) no. of moles of solute per liter of solution.

moles of soluteM = liter of solution

Example: What is the molar concentration of a solution containing 16.0g CH3OC in 200 mL of solution? Answer:


Chemistry

16.0g32.0g/mole

M=0.2 litermoleM=2.5liter

4. Molarity (m) no. of moles of solute per kg. of solvent.

number of moles solutem = kg solvent

Example: The molarity of a solution of C2H5OH in water is 1.25 mol/kg. How many grams of alcohol are dissolved in 2.5 kg of water? Answer:

1.25mole 46gg=2.5kg

kg moleg=143.75g

5. Normality (N) no. of equivalent weight of solute per liter of solution.

grams of soluteN = (eqv wt. of solute) x (liter of soln)

The equivalent weight of solute is determined by its change in valence in the particular reaction used. It follows that:

molecular mass (g/mole)eqv. wt. (g/eqv) = change in valence (eqv/mole)

Colligative Properties of Solutions

These are properties of solution which depend on the number of solute particles but not on the identity of the solute.

1. Vapor Pressure Lowering the addition of a non-volatile solute lowers the vapor pressure of the liquid because the solute reduces the fraction of solvent present. With relatively fewer solvent molecules, the rate of their escape from solutions is diminished, resulting in a decreased vapor pressure.

2. Boling Point Elevation addition of a non-volatile solute lowers the vapor pressure of the

solution. As a result, the boiling point of the solution will be higher than that of the pure liquid. 3. Freezing Point Depression the decrease in freezing point is directly proportional to the

molarity of the solute.


Chemistry

4. Osmotic Pressure pressure needed to prevent osmosis. (Osmosis the net movement of solvent molecules through a semi-permeable membrane from a more dilute solution to a more concentrated one.)

X. CHEMICAL KINETICS

Chemical Kinetics study of rates of chemical reactions, rate laws and reaction mechanisms. Reaction Rate number of moles of a reactant consumed per unit time. Usually the unit used is

molars per second (M/sec). Rate Law an equation or mathematical expression showing the relationship between reactant

concentrations and rate of reaction. Rate Constant an experimentally determined constant of proportionality between the reaction rate

and the concentrations of reactants that appear in the rate law. Law of Mass Action at constant temperature, the rate of reaction is usually proportional to some

power of concentration of each reactant. Order of Reaction the sum of the powers of the concentration factors in the rate equation. Reaction Mechanisms series of successive elementary steps by which reactants are converted

to products. Factors Affecting Reaction Rate:

Factors Effect on reaction rate

1. greater frequency of collision increase

2. higher energy of activation decrease

3. higher energy of activation Increase

4. lower temperature decrease

5. increasing the concentration of reactants Increase

6. increasing the particle size of reacting molecules decrease

7. using a catalyst Increase

Catalyst a substance that changes the rate of a chemical reaction without itself undergoing a

permanent change.

XI. CHEMICAL EQUILIBRIUM

Chemical Equilibrium a state in which two opposing chemicals reactions are proceeding in opposite directions at the same speed.

Le Chateliers Principle if a system at chemical equilibrium is disturbed by some stress, the system goes to a new equilibrium condition in such a way as to relieve the stress.

XII. ACIDS AND BASES

Properties of Acids and Bases


Chemistry

Acids Bases

Taste sour Turn blue litmus red Electrolytes React with metal to produce

hydrogen gas React with carbonates and

bicarbonates to produce CO2 gas Turn colorless with phenolphthalein

solution Turn red with methyl orange indicator pH values less than 7

taste bitter turn red litmus blue feel slippery electrolytes turn pink to violet color with phenolphthalein

solution turn yellow with methyl orange indicator pH values greater than 7

Conceptual Definitions of Acids and Bases Strong Electrolytes all ionic compounds that are dissolved in water (with very few exceptions) Weak electrolytes compounds with non-metallic cations and anions; their degree of ionization is

lower than those of strong electrolytes. Non-electrolytes compounds whose solution in water does not conduct electricity. Neutralization reaction between an acid and a base that produces a salt, a neutral compound. Hydrolysis a reaction between an ion and water. Amphoterism property of some compounds to behave both as an acid and as a base.

XIII. ORGANIC CHEMISTRY Definition Organic chemistry is the chemistry of the compounds of carbon with the exceptions of

carbon monoxide, carbon dioxide, carbonates group, and the cyanide group. Even though these compounds contain carbon, they were obtained from minerals and are considered to be inorganic compounds.

General Comparison of Organic and Inorganic Compounds

Property Organic Compounds Inorganic Compounds

1. Solubility in water Insoluble, except those that are capable of H-bonding Soluble

2. Solubility in organic solvents Soluble Insoluble

3. Melting point Low Very high

4. Boiling point Low Very high

5. Electrical conductivity Non-conductors Conduct in molten state of solution

6. Molecular mass High Low

7. Structure Complex Simple

8. Particles Molecules Ions

9. Combustion Mostly flammable Usually


Chemistry

10. Isomerism Common Rare


Chemistry

CRITICAL THINKING QUESTIONS

1. The temperature at which mercury starts to freeze is -35C. What is the temperature, in degrees

Fahrenheit, at which a mercury thermometer can not be used? a. -63 F b. -35 F c. -31 F d. -5.4 F 2. A piece of stone weighs 0.05 pounds. When if is submerged in a graduated cylinder containing 50 mL of

H2O, the level rose to 60 mL. What is the density of the stone in g/mL? a. 2.27 b. 2.72 c. 7.22 d. 7.27 3. A swimming pool 25m wide, 100m long, and 3m deep is filled with water up to a height of 2m. How many

kilograms of water have been placed inside the pool? (The density of water is 1000 kg/m3.) a. 5 million b. 7.5 million c. 10 million d. 12.5 million 4. Which of the following not an example of a compound? a. sugar b. salt c. ash d. water 5. If two pure substances have different melting points, then

a. the two substances will surely have different densities b. the two samples are certainly different pure substances c. the two substances are certain to have identical chemical formulas d. the two substances are certain to be compounds and not elements

6. Which of the following is a compound? a. water b. wine c. soil d. mercury 7. Which of the following processes is an example of a chemical change? a. evaporation of sea water to form salt b. melting of an ice cube

c. filtering of paper pulp from a liquid slurry using a sieve tray d. rusting of iron 8. We know that air is a homogeneous mixture and not a compound because a. it has no definite shape

b. it has no definite volume c. it ca be compressed d. its composition can vary 9. What do you call a substance that is composed of two or more elements bonded chemically? a. an isotope b. an element c. a compound d. a mixture 10. Which of the following examples is a physical change?

a. crystallization of sugar from sugar can juice b. fermenting of ethanol to form wine c. burning of a piece of candle d. clotting of blood

11. Which of the following substances cannot be further decomposed by ordinary chemical means? a. water b. sugar c. air d. silver 12. Which of the following not a manifestation of a chemical change?

a. reaction of a compound and an element to form a new compound and an element


Chemistry

b. breaking down of compound into elements c. combining of atoms of elements to form a molecule d. separation of the molecules in a mixture

13. What do you call a nuclear reaction resulting from the interaction of two nuclei to form a bigger nucleus

and an accompanying release of energy? a. nuclear fission b. alpha emission c. nuclear fusion d. natural radioactive decay 14. Which of the following materials cannot be subjected to carbon dating to determine its age? a. a trunk of wood b. a sword c. a smear of blood on a piece of cloth d. an ancient Egyptian scroll 15. What law states that the pressure of a gas is directly proportional to its absolute temperature at constant

volume? a. Charles Law b. Gay-Lussacs Law c. Boyles Law d. Daltons Law 16. To what conditions does a gas behave like an ideal gas? a. low temperature and low pressure b. low temperature and high pressure c. high temperature and low pressure d. high temperature and high pressure 17. What law states that the pressure of gas is inversely proportional to its volume at constant temperature? a. Charles Law b. Gay-Lussacs Law c. Boyles Law d. Daltons Law 18. Which gas diffuses faster? a. CH4 b. O2 c. CO d. He 19. Which of the following best describes heat? a. the capacity to do work b. forces times distance c. sum of thermal and chemical energy

d. an energy transfer due to a temperature difference 20. What happens to water when it begins to vaporize? a. it increases in temperature b. it decreases in temperature c. no change in temperature d. no change in thermal energy 21. Burning of gasoline initially requires heat before it burns spontaneously. Which of the following does not

give a good explanation of this phenomenon? a. the initial heat rises the enthalpy of the reactant b. the initial heat lowers the activation of energy of the reactants c. the enthalpy of the reactants is lower that the enthalpy of the products d. the enthalpy of the product is lower than the enthalpy of the reactant.


Chemistry

22. Which of the following is an endothermic process?

a. melting of ice b. burning of paper c. neutralization of a strong acid and a strong base d. violent reaction of sodium metal with water

23. Which of the following events is heat exchange involved? a. when there is a phase change b. when there is a chemical reaction c. when the gas expands adiabatically d. when there is difference in temperature 24. Who first predicted the wave-particle dual property of electrons? a. Hund b. Heisenberg c. De Broglie d. Schrdinger 25. Who postulated the wave equation that describes the properties of electrons in an atom? a. Bohr b. Heisenberg c. Pauli d. Schrdinger 26. Atoms of nonmetals generally reacts with atoms of metals by

a. gaining electrons to form ionic compounds b. gaining electrons to form covalent compounds c. sharing electrons to form ionic compounds d. sharing electrons to form covalent compounds

27. The addition of a nonvolatile solute to a solvent will cause

a. the vapor pressure of the solvent to increase b. the vapor pressure of the solvent to decrease c. the vapor pressure of the solvent to remain unchanged d. non of these

28. Which of the following factors does not affect the rate of reaction? a. the number of products formed b. the nature of reactants c. temperature d. concentration of reactants 29. Which of the following statements about the catalyst is not true?

a. they may slow down the reaction b. they may speed up the reaction c. they are present in living substances d. the may become new substances after the reaction

30. Which of the following statements about equilibrium is TRUE?

a. it exists in a closed system at varying temperature b. it exist in an open system c. it exists between a liquid and its vapor in a closed system at uniform temperature. d. it may exist between a solid and a liquid.

31. According to the Bronsted-Lowry theory, an acid is

a. a proton donor b. a proton acceptor c. a proton donor and a proton acceptor d. neither a proton donor nor a proton acceptor


Chemistry

32. Which is not true of bases? a. they always contain OH- ions b. they neutralize acid c. the pH of their solution is greater than 7 d. they react with H3O+ ions 33. Organic chemistry is the chemistry of compounds containing the element a. hydrogen b. carbon c. oxygen d. nitrogen 34. What is the mass in grams of 1 liter of carbon monoxide (CO) at standard temperature and pressure

(STP)? Note: The molecular weight (MW) of CO is 28 g/mole, and at STP, 1 mole of any gas occupies a volume of 22.4 liters.

a. 1.20 b. 1.35 c. 1.45 d. 1.25 35. Two-thirds of the atom in a molecule of water is hydrogen. What percentage weight of a water molecule

if the weight of two hydrogen atoms? The atomic weight of hydrogen is 1.008 g/mol and oxygen is 16.00 g/mole.

a. 19.12 b. 11.19 c. 19.11 d. 12.19 36. How many protons (P) and neutrons are there in the nucleus are present in a Pb nucleus of atomic

mass of 206? a. P = 92, N = 156 b. P = 85, N = 160 c. P = 82, N = 124 d. P = 90, N = 150 37. A 0.064 kg. of octane vapor (MW = 114) is mixed with 0.91 kg of air (MW =29.0) in the manifold of an

engine. The total pressure in the manifold is 86.1 kPa, and a temperature is 290 K. Assume octane behaves ideally. What is the partial pressure of the air in the mixture in KPa?

a. 46.8 b. 48.6 c. 84.6 d. 64.8 38. Hydrogen peroxide solution for hair bleaching is usually prepared mixing 5 grams of hydrogen peroxide

(H2O2), Molecular weight = 34 g/mole) per 100 ml of solution. What is the molarity of this solution? a. 1.0 M b. 1.5 M c. 1.95 M d. 1.8 M 39. A cylinder contains oxygen at a pressure of 10 atm and a temperature of 300K. The volume of the

cylinder is 10 liters. What is the mass of oxygen in grams? Molecular weight (MW) of oxygen is 32 g/mole?

a. 125.02 b. 130.08 c. 135.05 d. 120.04 40. The molecular diameter of CO is 3.19x10-8 at 300K and pressure of 100 mmHg. What is the mean free

path of the gas in cm? a. 6.86x10-3 b. 6.86x10-5 c. 2.86x10-4 d. 6.86x10-9 41. How many moles are there in one atom? a. 3.6x10-23 b. 1.66x10-5 c. 2.86x10-4 d. 6.86x10-9


Chemistry

42. When 0.5g of liquid is completely evaporated and collected in liter manometer, the pressure is 0.25 atm and the temperature is 27C. Assume ideal gas behavior, find the molecular weight if the gas constant is 0.0821 L.atm/mole.K.

a. 49.2 g/mole b. 12.3 g/mole c. 2.2 g/mole d. 64.0 g/mole 43. If the atomic weight of magnesium is 24.3 g/mol, calculate how many magnesium atoms does 5g

represent? a. 1.24x1023 atoms b. 1.76x1023 atoms c. 3.44x1023 atoms d. 2.76x1023 atoms 44. How many moles of iron does 25 g of Fe represent? Note: the atomic weigh of iron (Fe) is 55.8 g/mol. a. 0.356 mol b. 0.564 mol c. 0.448 mol d. 0.247 mol 45. How many oxygen atoms are present in 2.00 moles of oxygen molecules considering that is a diatomic? a. 2.40 x 1024 atoms b. 3.43 x 1025 atoms c. 5.67 x 1026 atoms d. 1.34 x 1024 atoms 46. if the atomic mass of copper (Cu) if 63.5 g/mol, compute how many grams does 0.252 mole of copper

(Cu) has? a. 16 g b. 18 g c. 20 g d. 12 g 47. What is the molecular weight of calcium hydroxide or Ca(OH)2? a. 74 b. 67 c. 80 d. 44 48. How many molecules are there in 25 g of hydrogen Chloride, HCl? a. 4.12 x 1023 molecules b. 4.32 x 1023 molecules c. 5.34 x 1023 molecules d. 3.45 x 1023 molecules 49. What is the percentage composition of a solution in the sodium chloride compound? a. 60.7% b. 34.6% c. 39.3% d. 50.7% 50. What is the composition of oxygen of potassium sulfate, K2SO4? a. 53.2% b. 36.7% c. 50.4% d. 43.4% 51. A 1.63 g of zinc when heated in air combined with 0.40 g of oxygen to form oxide of zinc. What is the

percentage composition of Zn in the compound formed? a. 80.3% b. 76.5% c. 19.7% d. 53.4% 52. Calculate how many moles of ammonia can be produced from 8 mol of hydrogen reacting with

nitrogen? a. 4.53 mol NH3 b. 7.76 mol NH3 c. 5.33 mol NH3 d. 4.57 mol NH3 53. How many molecules of water can be produced by reacting 0.010 mol of oxygen with hydrogen? a. 1.20 x 1022 molecules b. 1.32 x 1022 molecules


Chemistry

c. 2.34 x 1022 molecules d. 4.15 x 1022 molecules 54. If 2 liters of gas measured at STP weigh 3.23 g, what is the molecular weight of the gas? a. 36.2 g/mol b. 42.3 g/mol c. 24.7 g/mol d. 19.4 g/mol 55. An ethyl ether 691 mL weighs 1.65 g measured at 40C and 630 torr. Compute the molecular weight of

ethyl ether. a. 34.5 g/mol b. 43.5 g/mol c. 73.9 g/mol d. 67.5 g/mol 56. Calculate the specific gravity of Cl2 at STP. Note: the molecular weight of Cl2 is 71 g/mol. a. 3.45 b. 1.23 c. 2.46 d. 1.76 57. Compute the volume of oxygen at STP that can be formed from a 0.75 mole of potassium chlorate

(KClO3). a. 18.6 liters b. 16.8 liters c. 25.2 liters d. 23.2 liters 58. What pressure will be exerted by a 0.50 mol of gas in a 7 L container at 23C? a. 1.74 atm b. 2.05 atm c. 3.04 atm d. 1.32 atm 59. Compute how many moles of oxygen has are in a 70 L tank at 25C if the pressure is 2000 psi? a. 389.3 mol b. 453.4 mol c. 145.7 mol d. 247.4 mol 60. What is the molarity of the solution that contains 65 g of sucrose (C12H22O11) dissolved in 300 g of

water? a. 0.89 mole/kg b. 0.78 mole/kg c. 0.54 mole/kg d. 0.63 mole/kg 61. Calculate the number of moles of an ideal gas sample at 0.6 atmosphere and 87C occupies 0.450 liter. a. 0.0091 mole b. 0.0087 mole c. 0.0076 mole d. 0.0056 mole 62. One gram of hydrogen gas (H2) is combined with 10 g of helium (He) gas and confined at 20C and 5

atmospheres. What is the combined volume in liters? a. 14.4 liters b. 17.5 liters c. 16.4 liters d. 12.7 liters 63. What is the molarity of the solution if 150 g of KCl is dissolved in water to make 800 mL solution? a. 2.51 moles/L b. 2.25 moles/L c. 2.87 moles/L d. 1.53 moles/L 64. Compute how many grams of KCl must be dissolved in water so that it can produce a 400 L of 0.6 M

(molarity) solution? a. 17.904 g b. 14.281 g c. 11.541 g d. 12.653 g 65. What is the atomic weight of calcium if 2.25 g of pure calcium metal are converted to 3.13 g of pure

CaO? a. 49.8 g/mol b. 54.3 g/mol


Chemistry

c. 23.7 g/mol d. 40.9 g/mol 66. What is the equivalent weight of Sulfuric Acid? a. 49 b. 98 c. 23 d. 100 67. If 60 g of H2SO4 is dissolved in water to make a 1.5 L solution, find its normality N? a. 0.813 equiv/L b. 0.576 equiv/L c. 0.871 equiv/L d. 0.765 equiv/L 68. What is the equivalent weight of Mg(OH)2? a. 23 g/mol b. 29 g/mol c. 58 g/mol d. 20 g/mol 69. How many grams of H3PO4 are confined in 700 mL container if its normality is 0.5? a. 11.45 g b. 12.34 g c. 10.56 g d. 9.35 g 70. Which of the following is the simplest balanced equation of the given reaction? Na2CO3 + HCl NaCl +

H2O + CO2 a. Na2CO3 + 2HCl 2NaCl + H2O + CO2 b. Na2CO3 + 2HCl NaCl + 2H2O + CO2 c. 2Na2CO3 + HCl 2NaCl + H2O + CO2 d. 2Na2CO3 + HCl NaCl + H2O + 2CO2 71. If the molecular formula of water is H2O, then what is its molecular mass? a. 18 amu b. 19 amu c. 20 amu d. 1 amu 72. What is the molecular weight of barium chloride dehydrate (BaCl22H2O)? a. 234.4 amu b. 244.3 amu c. 270.5 amu d. 298.5 amu 73. Which of the following is the simplest balanced equation of the given Oxidation-Reduction Equation?

P + HNO3 + H2O NO + H3PO4 a. 2P + HNO3 + H2O NO + 2H3PO4 b. 3P + HNO3 + H2O NO + 3H3PO4 c. 3P + 5HNO3 + 2H2O 5NO + 3H3PO4 d. 3P + HNO3 + 2H2O 2NO + 3H3PO4 74. What type of bond results in form the sharing of electrons by two atoms? a. atomic bond b. covalent bond c. metallic bond d. ionic bond 75. Which of the following statements regarding organic substances is FALSE? a. Organic substances generally dissolve in high concentration acids b. All organic matter contains carbon c. Organic matter is generally stable at very high temperatures d. Organic substances generally do not dissolve in water 76. What do you call a substance that dissociates in solutions to produce positive and negative ions? a. base b. acid c. electrolyte d. solute 77. During a complete or partial neutralization of acids, what is the ionic compound formed? a. salt b. sugar c. potassium d. sulfur


Chemistry

78. Which of the following is most likely to prove that a substance is inorganic? a. the substance evaporates in room temperature and pressure b. the substance is heated together with copper oxide and the resulting gases are found to have no effect

on limestone c. analysis shows that the substance contains hydrogen d. the substance floats in water 79. Which of the following are found in the nucleus of an atom? a. electrons and protons b. electrons and neutrons c. protons and neutrons d. electrons, protons and neutrons 80. Which of the following elements and compounds is unstable in its pure form? a. Helium (He) b. Neon (Ne) c. Carbon dioxide (CO2) d. Sodium (Na) 81. What refers to the negatively charge component of an atom? a. electron b. proton c. neutron d. ion 82. Which of the following is the simplest type of reaction where two elements of compounds combine

directly to form a compound? a. directly combination or synthesis b. decomposition or analysis c. single displacement d. double displacement 83. What do you call the bonding that occurs in inert gases and other elements with full shells, primarily due

to attraction between dipole structures? a. ionic b. metallic c. covalent d. van de waals 84. If the heat of a solution is negative, heat is given off when the solute dissolves in the solvent. What type

of reaction is this? a. exothermic b. ideal c. endothermic d. efflorescent 85. What do you call materials that do not conduct electric current? a. conductor b. insulator c. semi conductor d. intrinsic material 86. What element is known as the lightest metal? a. aluminum b. manganese c. Magnesium d. Lithium 87. Which of the following is energy removal being applied? a. evaporation of water b. changing water to steam c. changing water to ice d. all of these 88. Halogens fall under what group in the periodic table? a. group VIA b. group VA


Chemistry

c. group IVA d. group VIIA 89. Which of the following is added to the drinking water distribution system for disinfection? a. Soda ash b. Chlorine c. Lime d. Iodine 90. What refers to the number of gram equivalent weights of solute per liter of solution? a. molarity (M) b. normality c. molarity (m) d. formality 91. Is the attraction between like molecules. a. absorption b. diffusion c. adhesion d. cohesion 92. What do you call a substance that cannot be decomposed any further by a chemical reaction? a. ion b. element c. molecule d. atom 93. One of the following is the standard pressure and temperature. Which on? a. 0C and zone atmosphere pressure b. 0C and zero pressure c. 0F and one atmosphere d. 32F and zero pressure 94. Which of the following is the strongest type of bonds? a. Van de Waals b. Metallic c. Covalent d. Ionic 95. Sublimation is a direct change from: a. solid to liquid phase b. solid to gaseous phase c. liquid to gaseous phase d. gaseous phase to liquid phase 96. What do you call hydrocarbons containing carbon to carbon double bonds? a. Alkanes b. Alkenes c. Alkynes d. None of these 97. How are materials containing atoms with less than valence electrons classified? a. an insulator b. a semi-conductor c. a conductor d. a compound 98. Which of the following has the characteristics of both metals and non-metals? a. conductors b. insulators c. metalloids d. meteors 99. Which are oxidizing and reducing agents in the following reactions? 2CCl4 + K2CrO4 2Cl2CO + CrO2Cl2 + 2 KCl a. there are no oxidizing and reducing agents in this reaction b. oxidizing agent: chromium; reducing agent: chlorine c. oxidizing agent: chlorine; reducing agent: carbon d. oxidizing agent: oxygen; reducing agent: chlorine 100. How are elements numbered 58 to 71 in the periodic table called? a. Lanthanons b. Actinons c. Earth metals d. Noble gas


Chemistry

101. What type of bonding in which electrostatic attraction is predominant? a. Ionic bonding b. Metallic bonding c. Covalent bonding d. Van der Waals bonding 102. What term refers to the passage of an electric current trough an electrolyte caused by an external

voltage source? Which one? a. electrolysis b. electromechanical action c. electrolyte d. piezoelectric effect 103. When all of the atoms of a molecule are the same, the substance is called _____. a. a compound b. a chemical c. an element d. an ion 104. Which of the following refers to the measure of the amount of negative ions in the water? a. acidity b. alkalinity c. turbidity d. molarity 105. What device produces electrical current by way of an oxidation-reduction reaction? a. generator b. galvanic cell c. metallic friction d. all of these 106. What is the maximum number of electrons that can be accommodated in the valence shell of an atom? a. 6 b. 8 c. 10 d. 12 107. Reactions generally proceed faster at high temperatures because of which of the following? a. the molecules are less energetic b. the molecules collide more frequently c. the activation energy is less d. the molecules collide more frequently and the activation energy is less 108. Which component of an atom has no electric charge? a. proton b. neutron c. coulomb d. electron 109. Adding more solute to an already saturated solution will cause the excess solute to settle to the bottom

of the container. What is this process called? a. precipitation b. hydration c. dehydration d. saturation 110. What is formed when acids will react with active metals? a. sulfur b. oxygen c. hydrogen d. chloride 111. How much is the pH content of an acid? a. between 4 and 6 b. between 2 and 7 c. between 1 and 5 d. between 0 and 7 112. The condition of a liquid electrolyte is measured in terms of its: a. specific gravity b. acid content c. voltage output d. current value


Chemistry

113. What is a substance that speeds up a chemical reaction without itself undergoing a chemical change? a. ingredients b. reactants c. solvent d. catalyst 114. How are elements numbered 90 to 103 in the periodic table called? a. alkali b. actinons c. earth metals d. tr ansition elements 115. What is defined as a value equal to the number of gram moles of solute per 1000 grams of solvent? a. Molarity (m) b. Normality c. Molarity (M) d. Formality 116. Which of the following is NOT a part of an atom? a. electron b. proton c. neutron d. coulomb 117. An element maybe defines as a substance, all atoms of which have the same: a. number of neutrons b. radioactivity c. atomic weight d. atomic number 118. How does all B families and group VII in the periodic table named? a. light metals b. rare earth metals c. non-metals d. transition metals 119. The device which measures the acid content of the cell is called _____. a. acid meter b. hydrometer c. hygrometer d. pyrometer 120. The vertical columns in the periodic table are called: a. Groups b. Sections c. Batches d. Families 121. What term is used to refer to a negatively charged ion? a. Anion b. Cathode c. Anode d. Cation 122. In a copper atom, the valence ring contains: a. no electrons b. one electrons c. two electrons d. four electrons 123. The elements along the dark line in the periodic table are referred to as _____. a. Light metals b. Metalloids c. Non-metals d. Heavy metals 124. What do you call an atom that loses some of its electron or accepts extra electrons from another

atom? a. Intrinsic b. Mole c. Neutron d. Ion 125. A _____ is a cell designed to produce electric current and can be recharged. a. Secondary cell b. electrolytic cell c. chemical cell d. battery 126. Which of the following groups in the periodic table the most strongly electronegative elements? a. Group IV b. Group V c. Group VIIA d. Group VIA


Chemistry

127. Which of the following statement is FALSE? a. In general, as reaction products are formed, they react with each other and reform reactants. b. At equilibrium, the net reaction rate is zero. c. The differential rate is the mathematical expression that shows how the rate of reaction depends on

volume. d. The net rate at which a reaction proceeds from left to right is equal to the forward rate minus the reverse

rate. 128. What is the smallest part of matter? a. molecule b. element c. particle d. atom 129. The opposite of alkali a. acid b. fluid c. substance d. none of these 130. What type of reaction has two compounds as reactants and two compounds as products? a. Direct combination or synthesis b. Decomposition or synthesis c. Single displacement d. Double displacement 131. If the heat of a solution is positive, heat is absorbed when the solute dissolves in the solvent. What

type of reaction is this? a. Exothermic b. Ideal c. Endothermic d. Efflorescent 132. The amount of electricity a battery can produce is controlled by a. the thickness of the plate b. the plate surface area c. the strength of the acid d. the discharge load 133. What do you call the electrons in the last orbit or shell of an atom? a. Bound electrons b. Free electrons c. Valence electrons d. External electrons 134. Which one is the positively charged ion? a. anion b. cathode c. anode d. cation 135. It is the number of protons in the nucleus of an atom a. Molecular number b. Proton number c. Mass number d. Atomic number 136. What do you call the elements in the first two groups in the periodic table? a. Light metals b. Noble gas c. Non-metals d. Heavy metals 137. Which one refers to the number of gram-moles of solute per liter of solution? a. Molarity (m) b. Normality c. Molarity (M) d. Formality 138. What type of bonding is electrostatic attraction most predominant? a. Ionic b. Metallic


Chemistry

c. Covalent d. Van der Waals 139. When the charge of an atom becomes unbalanced, the charge atom is called _____. a. an ion b. a neutron c. a proton d. an electron 140. Which of the following type of reactions in which bonds within a compound are disrupted by heat or

other energy to produce simpler compounds or elements? a. direct combination or synthesis b. decomposition or analysis c. single displacement d. double displacement 141. What do you call hydrocarbons containing carbon to carbon triple bonds? a. alkanes b. alkenes c. alkynes d. non of these 142. If an atom contains more than four valence electrons, the material is classified as _____. a. insulator b. semi-conductor c. conductor d. any of these 143. Which of the following refers to a measure of the quantity of an element or compound? a. Oxidation number b. Atomic number c. Avogadros number d. Mole 144. The electrolyte is a solution of water and _____. a. sulfuric acid b. uric acid c. nitric acid d. formic acid 145. Acids will turn blue litmus paper to what color? a. gray b. yellow c. violet d. red 146. How much is the pH content of a base? a. between 6 and 10 b. between 2 and 7 c. between 7 and 14 d. between 10 and 16 147. Applying a greater pressure causes pure solvent to leave the solution. What is the name of this

process? a. cavitation b. calcination c. purification d. reverse osmosis 148. A deuteron is a. a neutron plus two protons b. a nucleus containing a neutron and a proton c. an electron with a positive change d. a helium nucleus 149. Which of the following groups in the periodic table the most weakly electronegative elements? a. Group IIA b. Group IA c. Group IIIA d. Group IVA 150. What type of bonding that occurs in metals when metal atoms lose electrons and the metallic ions are

attracted to a sea of delocalized electrons? a. ionic bonding b. metallic bonding c. covalent bonding


Chemistry

d. van der waals bonding 151. Which of the following elements is NOT radioactive? a. Plutonium b. Californium c. Uranium d. Cobalt 152. Which of the following refers to a nucleic acid that stores genetic information? a. Cellulose b. Codon c. DNA d. Buffer 153. During chemical reactions, bonds between atoms are broken and new bonds are usually formed. What

do you call the starting substances? a. Products b. Reactants c. Catalyst d. Ingredients 154. What do you call solutions having the same osmotic pressure? a. Isotonic solutions b. Monohydroxic solutions c. Dihydroxic solutions d. Toxic solutions 155. What is the term used to describe hydrocarbons containing single covalent bonds between carbon

atoms? a. alkanes b. alkenes c. alkynes d. allotrope 156. The smallest whole unit of an element like Uranium is: a. molecule b. atom c. ion d. electron 157. What refers to the formation and collapse of minute bubbles of vapor in liquid which caused by a

combination of reduced pressure and increased velocity in the fluid? a. cavitation b. stress corrosion c. fatigue corrosion d. precipitation 158. The elements in Group 0 in the periodic table is called_____. a. light metals b. rare earth metals c. noble gas d. heavy metals 159. A pair of electrical conductors of dissimilar materials so joined as to produce a thermal emf when the

junctions are of different temperatures. a. potentiometer b. piezoelectric c. thermocouple d. solar heating 160. Bases will turn red litmus paper to what color? a. blue b. yellow c. violet d. green 161. Which of the following refers to the diffusion of a solvent into a stronger solution in an attempt to

equalize the two concentrations? a. purification b. electrolysis c. osmosis d. hydrolysis 162. The formula for Dinitrogen Pentoxide is: a. N2O5 b. (NO)5 c. NO d. none of these 163. Which of the following will occur if a substance is oxidized? a. it absorbs energy


Chemistry

b. it loses electrons c. it becomes more negative d. it gives off heat 164. What type of reaction has one element and one compound as reactants? a. direct combination or synthesis b. decomposition of analysis c. single displacement d. double displacement 165. Which of the following is the atomic number of silicon? a. 32 b. 24 c. 14 d. 28 166. Dielectric is another name for _____. a. a conductor b. an element c. an insulator d. a capacitor 167. Which of the following refers to the change from gaseous to liquid phase? a. condensation b. vaporization c. sublimation d. ionization 168. At the same, pressure and temperature, equal volumes of all gasses contain equal number of

molecules. This is known as a. Boyles Law b. Faradays Law c. Avogadros Law d. Charles Law 169. The galvanic cell is not dependent of which factor? a. temperature b. pressure c. volume d. chemical substance 170. Which of the following refers to the number of protons and neutrons in the nucleus of an atm? a. atomic weight b. atomic mass c. atomic constant d. atomic number 171. One of the following statements is wrong. Which one is it? a. Electron is an elementary quantity of negative electricity b. proton is an elementary quantity of positive electricity c. an atom is composed of a central nucleus and orbital electrons d. the mass of an electron is heavier than that of a proton 172. During a chemical reactions, bonds between atoms are broken and new bonds are usually formed.

The ending substance is called _____. a. products b. reactants c. catalyst d. ingredients 173. What is the smallest subdivision of an element of compound that can exist in a natural state? a. atom b. molecule c. ion d. element 174. The sharing of one or more electron pairs between nuclei. It usually occurs when the electronegativity

difference between bonding species is less than 1.5 a. Bridge Bonding B. Ionic Bonding C. Covalent Bonding D. Valence Bonding 175. The name of this group is the chalcogen (oxygen) family. a. Group IIA b. Group IVA c. Group VIA d. Group VIIIA


Chemistry

176. The amount of energy to change 1 g of liquid to gas at its boiling point. a. enthalpy of formation b. enthalpy of fusion c. enthalpy of reaction d. enthalpy of vaporization 177. The fragmentation of a crystal along a characteristic crystallographic direction, caused by lines of

weakness in constituent atomic groups. a. cleavage b. fracture c. luster d. streak 178. A bond formed by the sideways overlap of two parallel p orbitals. a. Peptide Bond b. Pi Bond c. Saturated Bond d. Delta Bond 179. This indicator turns into colorless if the substance is acidic. a. Phenolphthalein b. Methyl orange c. Litmus paper d. Bromthymol blue 180. A chemical structure with definite formula for which there exists one or more distinct structures with the

sane formula. a. Isomer b. Radical c. Group d. Colloid 181. A substance that can be decomposed into 2 or more simpler substances by ordinary chemical means. a. Element b. Mixture c. Compound d. Solution 182. A theory which treats bonding as an over lapping of ligand orbitals with those of the central atom. a. Ligand Field Theory b. Crystal Fields Theory c. Chelate Effect d. Molecular Orbital Theory 183. Known as fools gold. a. Quartz b. Cinnabar c. Feldspar d. Pyrite 184. It is the hardest substance known a. Wad b. Graphite c. Diamond d. Adamantium 185. The splitting of water to form hydrogen and oxygen a. reduction b. electrolysis c. oxidation d. hydrolysis 186. Which is not an open chain hydrocarbon classification? a. alkenes b. alkanes c. cycloalkanes d. alkynes 187. The process of producing ions from neutral species. a. Fractional Distillation b. Hydration c. Recombination d. Ionization


Chemistry

188. A chemical bond with sausage roll shape formed by the sideways overlap of two d orbitals. a. Peptide Bond b. Pi Bond c. Saturated Bond d. Delta Bond 189. Multidentate ligands have equal probability of forming a coordination bond as do monodentate ions. a. Ligand Field Theory b. Paulings Rule c. Chelate Effect d. Molecular Orbital Theory 190 It has no definite composition whose members are composed of two or more substances, each

retaining its own identifying properties. a. Homogeneous Mixture b. Heterogeneous Mixture c. Aqueous Mixture d. Ingeneous Mixture 191. A chemical compound having one or more unpaired electrons which is capable of bonding with

another compound. a. Labile b. Resonance Hybrid c. Free Radical d. Isodemic Crystal 192. The mineral name for common table salt a. Halite b. Cassierite c. Aragonite d. Calcite 193. Any process that involves the loss of electrons. a. Reduction b. Exsolution c. Oxidation d. Hydrolysis 194. Elements normally found in combination with iron and nickel. a. Noble Metals b. Promethium c. Tritium d. Siderophile 195. Any substances that changes the rate of a reaction without being used up. a. Catalyst b. Enztme c. Reactor d. Stimulus 196. A/An ____ of chemical substances is, by definition, the number in grams corresponding to the atomic

or molecular mass. a. atiomic mass b. atomic weight c. mole d. ion 197. A reaction in which an unstable reaction intermediate is found a. Coupled Substitution b. Free Radical Reaction c. Polar Reaction d. Concentrated Reaction 198. A reaction for which the difference between enthalpies of formation between products and reactants is

positive. a. Intrathermic b. Exothermic c. Endothermic d. Isothermic


Chemistry

199. The _____ is the heat released of absorbed in a chemical reaction at constant pressure when simple substances combine into a more complex substance.

a. enthalpy of formation b. enthalpy of fusion c. enthalpy of reaction d. enthalpy of vaporization 200. Characteristic based upon the reaction of a substance with other materials. a. Chemical Property b. Colligative Property c. Nuclear Property d. Physical Property 201. A symmetrical intergrowth of two or more crystals of the same substance. a. Twin b. Isomer c. Epitaxis d. Habit 202. A compound which is gaseous at ambient temperature and pressure and which is easily melted. A. Volatile b. Soluble c. Liquidus d. Solidus 203 The energy required to form gaseous monatomic species. a. Bond Energy b. Dissociation Energy c. Activation Energy d. Light Energy 204. The amount of energy released as on mole of a given substance is burned in the presence of oxygen. a. Nuclear Energy b. Enthalpy of Reaction c. Trans Effect d. Heat of Combustion 205. The most diamagnetic naturally occurring material. a. Iron b. Gallium c. Bismuth d. Silicon 206. Any Process that involves the gain of electrons. a. Reduction b. Exsolution c. Oxidation d. Hydrolysis 207. The change of the thermodynamic state function enthalpy due to a chemical reaction. a. enthalpy of formation b. enthalpy of fusion c. enthalpy of reaction d. enthalpy of vaporization 208. A concept invented by Linus Pauling to measure the tendency for atoms to form ionic instead of

covalent bonds. a. Electromagnetism b. Electropositivity c. Electronegativity d. Electrodynamism 209. A substance whose particles are strongly attracted to each other (e.g., gelatin) a. Colligative b. Hydrophilic c. Hydrophobic d. Cohesive 210. The most common form of calcium carbonate. a. Halite b. Cassierite c. Aragonite d. Calcite


Chemistry

211. The substance that undergoes phase change in the process of dissolving a. Radical b. Isomer c. Solvent d. Solute 212. A process in which water molecules are attracted to and form weak bonds with the solute species. a. Hydration b. Oxidation c. Combustion d. Reduction 213. The dissociation of a chemical species resulting from its absorption of a photon. a. Photosynthesis b. Photolithography c. Photosensitization d. Photodissociation 214. _____ is the mineral name for lead sulfide (PbS). It is the most important ore of lead. a. Saltpeter b. Gypsum c. Galena d. Silicate 215. A substance that does not undergo chemical reactions is said to be _____. a. inert b. labile c. alkaline d. amorphous 216. Which bond strength is weak? a. Macromolecular covalent bonds b. Ionic bonds c, Hydrogen bonds d. Van der Waals forces 217. Occurs with the bonding electron pair remaining intact. a. Coupled Substitution b. Free Radical Reaction c. Polar Reaction d. Concentrated Reaction 218. A mixture composed of a suspended and a mobile phase. a. Isomer b. Radical c. Suspension d. Colloid 219. A/An _____ is a biological catalyst. a. Amino acids b. Genes c. DNA d. Enzyme 220. The first covalent bond formed between two nuclei is always a _____. They are formed when two s

orbitals, one s and one p orbital, two p orbitals, or two d orbitals overlap. a. Peptide Bond b. Pi Bond c. Sigma Bond d. Bridge Bond 221. The process by which solutions are decomposed into their components by using differences in their

boiling points. a. Fractional Distillation b. Filtration c. Recombination d. Ionization 222. Filtration through a semi-permeable membrane used to separate colloids. a. Dialysis b. Hydrolysis c. Electrolysis d. Chromatography


Chemistry

223. In orbitals of identical energy, electrons remain unpaired if possible in order to minimize electron-electron repulsion.

a. Spin Multiplicity Rule b. Selection Rules c. Hunds Rule d. Paulings Rules 224. Elements with the same outermost shell are said to belong to the same _____. a. spin b. period c. valence d. slot 225. An alkyl sulfate alkali. a. Detergent b. Acetone c. Ethylene d. Oil 226. The process whereby an excited species transfers its energy to another excited species which

subsequently undergoes a reaction a. Photosynthesis b. Photolithography c. Photosensitization d. Photodissociation 227. _____ is the study of hydrocarbon compounds, ie., substances consisting of the elements hydrogen,

carbon, and oxygen. a. Inorganic Chemistry b. Organic Chemistry c. Quantum Chemistry d. Biochemistry 228. _____ displays all chemical elements systematically in order of increasing atomic number. a. Phase Diagram b. Polarization Spectrum c. Periodic Table d. Elements Archive 229. A substance composed of a single type of atom a. Atom b. Compound c. Solution d. Colloid 230. A group of elements which are gaseous at room temperature and pressure, and called so because

they rarely bond with other elements. a. Tetratomic Elements b. Nobel Gases c. Alkali d. Lanthanide 231. An element which is not found naturally on Earth. It has been found in the star HR465 in Andromeda. a. Adamantium b. Promethuim c. Tritium d. Adolinium 232. A form of elemental carbon which, because of its sheet structure, is an excellent lubricant. a. graphite b. Carbon c. Soot d. Coal 233. Which is not one of the periodic properties? a. electronegativity b. electron affinity c. radioactive decay coefficient d. atomic radius 234. The reflection of light beam passing through a colloid which identifies the presence of suspended

particles. a. Tyndall Effect b. Brownian Movement c. Raoults Law d. Trans Effect 235. An ionic compound containing a halogen.


Chemistry

a. Halide b. Lanthanide c. Alkalide d. Sulfide 236. A reaction for which the difference between enthalpies of information between products and reactants

is negative. a. Intrathermic b. Exothermic c. Endothermic d. Isothermic 237. The quantity of energy released as one mole of bonds are produced between atoms. a. Bond Energy b. Light Energy c. Activation Energy d. Atomization Energy 238. An ion with negative charge a. Cation b. Anion c. Muon d. Neutrino 239. _____ is a measure of the acidity of alkalinity of a substance. a. HP b. dB c. pH d. Pro-V 240. Two-phase mixture composed of a dispersed and continuous phase. a. Isomer b. Radical c. Suspension d. Colloid 241. A solvent for atoms. a. Flux b. Water c. Inert Gases d. Kryptonite 242. A class of matter with definite properties whose members are composed of two or more substances,

each retaining its own identifying properties. a. Homogeneous Mixture b. Heterogeneous Mixture c. Solid Solution d. Ingeneous Matter 243. A crystal for which all bonds have the same electrostatic valency. a. Labile b. Resonance Hybrid c. Free Radical d. Isodemic Crystal 244. The _____ is a system of reference materials against whose hardness a sample is compared. a. Richter scale b. Mohs hardness scale c. Cavendish balance d. Brinell hardness model 245. Magma with most of the gas component escaped. a. Lava b. Cinder c. Sulphur d. Coal 246. Minerals having high melting temperatures. a. Endogenic b. andesitic c. Granitic d. Basaltic 247. _____ minerals are those having low melting temperatures. a. Endogenic b. Andesitic c. Granitic d. Basaltic


Chemistry

248. The transfer of one or more electrons from a metal to a nonmetal. Electron transfer usually occurs when the electronegativity difference between bonding species is 1.5 or more.

a. Bridge Bonding b. Ionic Bonding c. Covalent Bonding d. Valence Bonding 249. The non-random overgrowth of two compositionally different crystalline substances. a. Twin b. Isomer c. Epitaxis d. Habit 250. The angles between equivalent faces of crystals of the same substance, measured at the same

temperature, are constant. a. Bravais Law b. Hall Effect c. Tyndall Effect d. Steno Law 251. The chemical element having the greatest binding energy per nucleon. a. Iron b. Gallium c. Bismuth d. Platinum 252. Any type of reaction that involves the pairing of unpairing of electrons. a. Coupled Substitution b. Free Radical Reaction c. Polar Reaction d. Concentrated Reaction 253. It is the form of oxygen consisting of three bound oxygen atoms. a. Halogen b. Halide c. Oxide d. Ozone 254. A group of elements with similar properties in which the outer-most electron shell is a partially filled f

sublevel. a. Tetratomic Elements b. Noble Gases c. Alkali d. Lanthanide 255. A substance whose particles are only weakly attracted to each other (e.g., water) a. Colligative b. Hydrophilic c. Hydrophobic d. Refractory 256. A class of homogeneous matter which has a definite composition by weight. a. compound b. mixture c. element d. substance 257. In the periodic table, Group IA is for the: a. Halogens b. Light metals c. Alkaline metals d. Alkali metals 258. A covalent bond formed through a condensation reaction that involves removal of a water molecule. a. Peptide Bond b. Pi Bond c. Sigma Bond d. Bridge Bond 259. Which group is the nitrogen family? a. Group VA b. Group IVA c. Group IIIA d. Group IIA 260. Litmus paper turns into _____ is the substance is basic. a. colorless b. yellow c. red d. blue


Chemistry

261. A property which depends only on the number of particles present, and not their chemical composition. a. Chemical Property b. Colligate Property c. Nuclear Property d. Physical Property 262. The amount of energy to change 1g of solid to liquid at its melting point. a. enthalpy of formation b. enthalpy of fusion c. enthalpy of reaction d. enthalpy of evaporation 263. The process whereby an initially homogeneous solid solution separates into two (or more) distinct

crystalline minerals without the addition or removal of materials to or rom the system. a. Reduction b. Exsolution c. Oxidation d. Hydrolysis 264. The separation of component particles from the bulk or mass. a. Dissociation b. Atomization c. Fission d. Reduction 265. An ion with positive change. a. Cation b. Anion c. Muon d. Neutrino 266. A cell which uses the flow of electrons from a spontaneous chemical reaction to do outside work. a. Daniel Cell b. Gravity Cell c. Concentration Cell d. Galvanic Cell 267. The anion OH-. a. Hydride b. Oxide c. Hydroxide d. Nitrate 268. A measure of the tendency of a gas to escape or expand. a. Compressibility b. Fugacity c. Vaporizability d. Solubility 269. An ionic theory which is an offshoot of electrostatic theory. It ignores all covalent bonding effects. a. Ligand Field Theory b. Crystal Fields Theory c. Chelate Effect d. Molecular Orbital Theory 270. A purple form of quartz whose color arises from Fe+4. a. Amethyst b. Ruby c. Topaz d. Peridot 271. A/An _____ is a hydrocarbon consisting only of single carbon-carbon bonds a. Alkyne b. Alkene c. Alane d. Amide 272. A large polypeptide of the kind found in living organisms. a. Amino Acid b. Protein c. Minerals d. Fatty Acid 273. It is the code in which almost all genetic information is encoded.


Chemistry

a. Transfer RNA b. Ribosomal RNA c. Ribonucleic Acid d. Deoxyribonucleic Acid 274. A sample of glass is a supercooled liquid rather than a true solid because it has _____. a. a definite volume b. no definite volume c. a crystalline structure d. no crystalline structure 275. Two immiscible liquids, when shaken together, may form a _______. a. solution b. sediment c. hydrated solution d. colloidal dispersion 276. Which of the following is not a pure substance? a. water b. milk c. hydrogen d. oxygen 277. Which of the following is not an isotope of hydrogen? a. hydrogen b. deuterium c. tritium d. uranium 278. Which of the following is formed by the transfer of electrons from one atom to another? a. allotrope b. ion c. isotope d. molecule 279. The modern periodic table is based on atomic _______. a. number b. radius c. charge d. mass 280. The gram molecular mass for H2SO4 is ______ grams. a. 7 b. 64 c. 98 d. 196 281. The rate of chemical reaction can be increased by _______. a. increasing the surface area of the reactants b. decreasing the reaction temperature c. decreasing the concentration of the reactants d. removing the catalyst 282. Which of the following is not a characteristic of a base? a. has a bitter taste b. feels smooth and slippery c. turns litmus from red to blue d. typically reacts vigorously with metals 283. Oxidation involves reactions in which ______. a. a substance gains one or more electrons b. a substance loses one or more electrons c. oxygen has been removed from the reactants d. oxygen is gained by the reactants 284. A bleach may work if it ______. a. adds hydrogen ions to the stain b. removes electrons during oxidation c. adds electrons to reduce coloration d. moves electrons between energy levels


Chemistry

285. An organic compound ______. a. contains carbon b. can be produced synthetically c. both 1 and 2 are correct d. neither 1 nor 2 is correct 286. Which of the following can be used to bombard atoms? a. protons b. neutrons c. electrons d. all three 287. An atom of which of the following elements has the greatest ability to attract electrons? a. silicon b. bromine c. sulfur d. nitrogen 288. Given the same conditions of temperature, which nobel gas will diffuse most rapidly? a. Kr b. He c. Ne d. Ar 289. When most fuels burn, the products include carbon dioxide and ______, a. hydrogen b. hydrocarbons c. water d. hydroxide 290. Which type of reaction is occurring when a metal undergoes corrosion? a. neutralization b. polymerization c. saponification d. oxidation-reduction 291. Which part of the Periodic Table contains elements with the strongest metallic properties? a. upper left b. lower left c. upper right d. lower right 292. A 1 M solution contains 20 grams of solute in 500 milliliters of solution. What is the mass of 1 mole of the solute? a. 10 g b. 20 g c. 40 g d. 80 g 293. What is the molecular formula of a compound with an empirical formula of CH and a molecular mass of 78? a. C6H6 b. C3H3 c. C4H4 d. C5H5 294. What is the empirical formula of a compound that contains 85% Ag and 15% F by mass? a. Ag2F b. AgF2 c. Ag2F2 d. AgF 295. A gas has a pressure of 300 torr, a temperature of 400 K, and a volume of 50.0 milliliters. What volume will the gas have at a pressure of 150 torr and a temperature of 200 K? a. 25.0 mL b. 50.0 mL c. 75 mL d. 100 mL 296. How many atoms are present in the formula KAl(SO4)2? a. 6 b. 8 c. 10 d. 12 297. How many liters of gas would 1.5 moles occupy at STP? a. 15.0 b. 4.5 c. 33.6 d. 44.6


Chemistry

298. If an element has an atomic number of 11, it will combine most readily with an element that has an atomic number of _____. a. 16 b. 17 c. 18 d. 19 299. Which of the following statements is true? a. chlorine changes in color b. chlorine changes in density c. chlorine changes into liquid d. chlorine reacts explosively to form sodium chloride 300. Which of the following is a compound? a. milk b. gold c. table salt d. ink 301. Golds atomic number is 79. How many protons, neutrons, and electrons are there in 197 Au? a. 79 b. 118 c. 197 d. 276 302. Which of the following is an element? a. H2O b. O3 c. CO2 d. C2H4 303. Which of the following is the empirical formula for glucose, a substance known as blood sugar, whose molecular formula is C6H12O6? a. CH2O b. CHO c. C2H4O d. CHO6 304. Which of the following is the name of the compound K2SO4? a. potassium sulfate b. sulphuric oxide c. potassium oxide d. silver nitrate 305. Which of the following balances this equation: Na(s) + H2O(l) NaOH(aq) + H2(g)? a. Na(s) + 2H2O(l) NaOH(aq) + H2(g) b. 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g) c. Na(s) + 2H2O(l) NaOH(aq) + H2(g) d. Na(s) + 2H2O(l) NaOH(aq) + H2(g) 306. What is the formula weight of C12H22O11(sucrose)? The atomic weights of carbon, hydrogen and oxygen are 12.0 amu, 1.0 amu and 16.0 amu, respectively. a. 29.0 amu b. 47.0 amu c. 76.0 amu d. 342.0 amu 307. What is the mass of 1 mol of glucose, C6H12O6? The atomic weights of carbon, hydrogen and oxygen are 12.0 amu, 1.0 amu and 16.0 amu, respectively. a. 180.0 amu b. 360.0 amu c. 53.0 amu d. 90.0 amu 308. What is the molarity of a solution made by dissolving a 23.4 g of sodium sulfate (NaSO4) in enough water to form 125 mL of solution? a. 2.12 M b. 1.32 M c. 3.33 M d. 5.34 M


Chemistry

309. What is the kinetic energy in joules and calories of a 6.0 kg object moving at a speed of 5.0 m/s? a. 30 cal b. 18 cal c. 1.2 cal d. 0.16 cal 310. What is the formula for a compound formed between aluminum and oxygen? (The atomic number of aluminum is 13 while that of oxygen is 8). a. Al2O3 b. Al3O5 c. AlO2 d. Al3O4 311. What is the volume of exactly 1 mol of gas at 0C (273.15 K) and exactly 1 atm pressure? a. 27.31 L b. 31.32 L c. 22.41 L d. 17.42 L 312. Which of the following substances is most likely to exist as a gas at room temperature and normal atmospheric pressures? a. P4O10 b. Cl2 c. AgCl d. I2 313. Which of the following will happen if a gas in an enclosed container is heated? a. pressure increases b. temperature decreases c. volume increases d. volume decreases 314. How can a bigger crystals of table salt be commercially produced? a. slow solar evaporation b. fast solar heating c. boiling in a cauldron d. heating over a sand bath 315. The natural fragrance of plants is attributed to the presence of ______. a. alkanes b. esters c. alcohols d. acetone 316. A solution is made containing 6.9 g of NaHCO3 per 100 g of water. What is the weight percentage of solute in this solution? a. 93.5% b. 89.7% c. 6.5% d. 10.3% 317. In terms of total mass, carbon monoxide (CO) is the most abundant of all pollutant gases. The most serious source of carbon monoxide poisoning comes from ______. a. cigarette smoking b. smoke from factories c. smoke from vehicles d. smoke from forest fires 318. Which of the following is not a non conventional source of biogas? a. animal manure b. petroleum c. ipil-ipil d. cassava 319. When mercury is placed in a small test tube, a

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