Topic Review Packet
Table of Contents
Topic 1: Matter, Its Properties & Changes Outline 3 PracticeQuestions . 5 Topic 2: Atomic Concepts Outline 8 Practice Questions.. 10 Topic 3: Periodic Table Outline . 17 Practice Questions .. 19Topic 4: Formulas & Names, Equations, Moles,
Molar Mass, & Types of Reactions Outline . 25 PracticeQuestions .. 26 Topic 5: Bonding Outline . 33 Practice Questions ..35 Topic 6 Overview .... 42 Topic 6A: Heat & TemperatureOutline . 43 Practice Questions .. 45 Topic 6B: Kinetics &Equilibrium Outline . 50 Practice Questions .. 51 Topic 7: Water& Solutions Outline .... 55 Practice Questions .. 56 Topic 8:Acids & Bases Outline . 61 Practice Questions .. 62 Topic 9:Organic Chemistry Outline . 66 Practice Questions .. 67 Topic 10:Phases & Gases Outline . 71 Practice Questions .. 72 Topic 11:Electrochemistry (Oxidation-Reduction) Outline . 75 PracticeQuestions .. 76 Answer Key 83
Topic 1: Matter, Its Properties and Changes Outline 1. Matter isclassified as a pure substance or a mixture of substances. A puresubstance (element or compound) has a constant composition andconstant properties throughout a given sample, and from sample tosample. You can use particle models/diagrams to differentiate amongelements,
compounds, and mixtures. 2. The proportions of components in amixture can be varied. Each component in a mixture retains itsoriginal properties. Differences in properties such as density,particle size, molecular polarity, boiling point and freezingpoint, and solubility permit physical separation of the componentsof the mixture. Methods of separating mixtures include evaporation,filtration, distillation,
and chromatography. Mixtures can be homogeneous orheterogeneous. Solutions are always
homogeneous. Heterogeneous mixtures are things like soil, fruitsalad, where the composition is NOT uniform throughout themixture.
3. The structure and arrangement of particles and theirinteractions determine the physical state of a substance at a giventemperature and pressure. Know the states (phases) of the elementsat STP; Br and Hg are the only
2 liquids, the noble gases as well as N, O, F, H, and Cl aregases, the rest are solids
Know the 7 elements that are diatomic in their natural states;7-Up or HOFBrINCl.
Draw particle models of solids, liquids, and gases. 4. Aphysical change results in the rearrangement of existing particlesin a substance; no new types of particles result from this type ofchange. A chemical change results in the formation of differentparticles with changed properties. Distinguish between chemical andphysical changes based on whether
new substances form or not.
5. Properties can be physical or chemical. Physical propertiesdescribe those characteristics that can be observed with the sensesor measured. Chemical properties describe how the substanceinteracts with other substances. Distinguish between chemical andphysical properties. One of the more useful properties is density.The density equation is on
Table T; D=m/V. Some common properties of the elements are foundon Table S, such as
melting and boiling points.
Matter Cut from Jan 2007 Jan 2008 Exams
1. A sample composed only of atoms having the same atomic numberis classified as
(1) a compound (3) an element (2) a solution (4) an isomer
2. A dilute, aqueous potassium nitrate solution is bestclassified as a (1) homogeneous compound (2) homogeneous mixture(3) heterogeneous compound (4) heterogeneous mixture 3. At whichCelsius temperature does lead change from a solid to a liquid?
(1) 874C (3) 328C (2) 601C (4) 0C
4. Which statement describes a chemical property of hydrogengas? (1) Hydrogen gas burns in air. (2) Hydrogen gas is colorless.(3) Hydrogen gas has a density of 0.000 09 g/cm3 at STP. (4)Hydrogen gas has a boiling point of 20. K at standard pressure. 5.Which element has the greatest density at STP?
(1) calcium (3) chlorine (2) carbon (4) copper
6. Which statement describes a chemical property of the elementmagnesium? (1) Magnesium is malleable. (2) Magnesium conductselectricity. (3) Magnesium reacts with an acid. (4) Magnesium has ahigh boiling point. 7. Matter that is composed of two or moredifferent elements chemically combined in a fixed proportion isclassified as
(1) a compound (3) a mixture (2) an isotope (4) a solution
8. Which element is a solid at STP and a good conductor ofelectricity?
(1) iodine (3) nickel (2) mercury (4) sulfur
9. The table below shows mass and volume data for four samplesof substances at 298 K and 1 atmosphere.
Which two samples could consist of the same substance?
(1) A and B (3) B and C (2) A and C (4) C and D
10. Bronze contains 90 to 95 percent copper and 5 to 10 percenttin. Because these percentages can vary, bronze is classifiedas
(1) a compound (3) a mixture (2) an element (4) a substance
11. At STP, which list of elements contains a solid, a liquid,and a gas?
(1) Hf, Hg, He (3) Ba, Br2, B (2) Cr, Cl2, C (4) Se, Sn, Sr
12. A 10.0-gram sample of which element has the smallest volumeat STP?
(1) aluminum (3) titanium (2) magnesium (4) zinc
13. At room temperature, a mixture of sand and water can beseparated by
(1) ionization (3) filtration (2) combustion (4) sublimation
14. Which particle diagram represents a sample of one compound,only?
15. A 1.00-mole sample of neon gas occupies a volume of 24.4liters at 298 K and 101.3 kilopascals. Calculate the density ofthis sample. Your response must include both a correct numericalsetup and the calculated result. Base your answers to questions 16through 18 on the information below.
In an investigation, a dripless wax candle is massed and thenlighted. As the candle burns, a small amount of liquid wax formsnear the flame. After 10 minutes, the candles flame is extinguishedand the candle is allowed to cool. The cooled candle is massed.
16. Identify one physical change that takes place in thisinvestigation. 17. State one observation that indicates a chemicalchange has occurred in this investigation.
18. Draw a particle diagram showing the change from solid wax toliquid wax. Use for particles of wax. Draw separate diagrams forthe liquid and the solid states.
Base your answers to questions 19 through 21 on the particlediagrams below, which show atoms and/ or molecules in threedifferent samples of matter at STP.
19. Which sample represents a pure substance? 20. When two atomsof y react with one atom of z, a compound forms. Using the numberof atoms shown in sample 2, what is the maximum number of moleculesof this compound that can be formed? 21. Explain why xx does notrepresent a compound.
Topic 2: Atomic Concepts Outline 1. The modern model of the atomhas evolved over a long period of time through the work of manyscientists. Daltons Model:
Elements are made of atoms Atoms of an element are the same.Compounds are formed from combinations of atoms.
Rutherford Experiment Bombarded gold foil with alpha particles.Showed atoms
were mostly empty space with small, dense positively chargednucleus.
Bohr Model Small, dense, positively charged nucleus surroundedby electrons in circular orbits.
Wave-Mechanical Model (Modern Atomic Theory) Small, dense,nucleus positively charged nucleus
surrounded by electrons moving in electron cloud. Orbitals areareas where an electron with a certain amount of energy is mostlikely
to be found. 2. Each atom is made of a positively chargednucleus with one or more orbiting, negatively charged electrons. 3.Protons and neutrons are found in the nucleus. The number ofprotons in an atoms nucleus gives the nucleus a positive charge. Lihas
a nuclear charge of +3, since it has 3 protons. 4. Protons havea positive charge, neutrons no charge, and electrons a negativecharge. 5. The number of protons in an atom equals the number ofelectrons. The positive charges of the protons are cancelled by thenegative charges of the electrons,
so overall an atom has a neutral charge. 6. The mass of a protonis 1 amu. The mass of a neutron is 1 amu. The mass of an electronis almost 0 amu. The mass of an atom is contained in its nucleus.The atomic mass of an atom is equal to the total number of protonsand neutrons. 7. Each electron in an atom has its own distinctamount of energy. When all electrons are at their lowest possibleenergy, it is called the ground state. Electrons fill in energylevels and orbitals starting with the one that requires the leastenergy
and progressively move to those levels and orbitals that requireincreasing amounts of energy.
8. When the electron gains a specific amount of energy, it movesto a higher orbital and is in the excited state.
You can recognize an excited state electron configuration. Ifthe configuration does not match that on the Periodic Table forthat number of electrons, then it is an excited state.
9. When an electron returns from a higher energy state to alower energy state, it emits a specific amount of energy usually inthe form of light. This can be used to identify an element (brightline spectrum). The instrument used to see the bright line spectrumis called a spectroscope. 10. The outermost electrons are calledvalence electrons. These affect the chemical properties of theelement. Atoms with a filled valence level are stable (noblegases). Most elements can have up to 8 electrons in their valencelevel. The exceptions are H and
He, which can have only 2 valence electrons. Atoms form bonds inorder to fill their valence levels. You can use Lewis structures toshow the configuration of the valence electrons. 11. Atoms of thesame element all contain the same number of protons. Changing thenumber of protons changes the atom into a different element. Theatomic number is the number of protons in an atom of an element.12. Isotopes are atoms with equal numbers of protons but differentnumbers of neutrons. Isotopes of an element have the same atomicnumber (protons only), but different atomic
masses (protons + neutrons). 13. The average atomic mass of anelement is the weighted average of its naturally occurringisotopes. You need to know how to do the calculation of weightedatomic mass given isotope
masses and percent abundances. 14. When an atom gains anelectron, it becomes a negative ion and its radius increases. 15.When an atom loses an electron, it becomes a positive ion and itsradius decreases. 16. Electronegativity indicates how strongly anatom of an element attracts electrons in a chemical bond. Thesevalues are based on an arbitrary scale. Fluorine has the highestelectronegativity of all elements (4.00).
Atomic Structure Practice Questions
1. Experiments performed to reveal the structure of atoms ledscientists to conclude that an atoms
(1) positive charge is evenly distributed throughout its volume(2) negative charge is mainly concentrated in its nucleus (3) massis evenly distributed throughout its volume (4) volume is mainlyunoccupied
2. The modern model of the atom shows that electrons are (1)orbiting the nucleus in fixed paths (2) found in regions calledorbitals (3) combined with neutrons in the nucleus (4) located in asolid sphere covering the nucleus
3. An experiment in which alpha particles were used to bombardthin sheets of gold foil led to the conclusion that an atom iscomposed mostly of
(1) empty space and has a small, negatively charged nucleus (2)empty space and has a small, positively charged nucleus (3) alarge, dense, positively charged nucleus (4) a large, dense,negatively charged nucleus
4. What is the atomic number of an element that has six protonsand eight neutrons? (1) 6 (2) 2 (3) 8 (4) 14
5. An atom of fluorine has a mass of 19 atomic mass units. Thetotal number of protons and neutrons in its nucleus is (1) 9 (2) 10(3) 19 (4) 28
6. What is the total number of protons contained in the nucleusof a carbon-14 atom? (1) 6 (2) 8 (3) 12 (4) 14
7. What is the nuclear charge of an iron atom?
(1) +26 (2) +30 (3) +56 (4) +82 8. Which of these elements hasan atom with the most stable outer electron configuration?
(1) Ne (2) Cl (3) Ca (4) Na 9. How many electrons are in theoutermost principal energy level of an atom of carbon in
the ground state? (1) 6 (2) 2 (3) 3 (4) 4
10. Which electron configuration is correct for a sodium ion?(1) 2-7 (2) 2-8 (3) 2-8-1 (4) 2-8-2
11. What is the electron configuration of a sulfur atom in theground state?
(1) 2-4 (2) 2-6 (3) 2-8-4 (4) 2-8-6
12. The nucleus of which atom contains 48 neutrons? (1) S3216(2) Ti
4822 (3) Rb
8537 (4) Cd
13. The number of neutrons in the nucleus of an atom can bedetermined by
(1) adding the atomic number to the mass number (2) subtractingthe atomic number from the mass number (3) adding the mass numberto the atomic mass (4) subtracting the mass number from the atomicnumber
14. When an atom loses an electron, the atom becomes an ion thatis
(1) positively charged and gains a small amount of mass (2)positively charged and loses a small amount of mass (3) negativelycharged and gains a small amount of mass (4) negatively charged andloses a small amount of mass
15. In which pair of elements do the nuclei of the atoms containthe same number of
neutrons? (1) Li73 and Be
94 (3) Na
2311 and Mg
(2) N147 and O168 (4) S
3216 and Cl
16. The characteristic spectral lines of elements are causedwhen electrons in an excited
atom move from (1) lower to higher energy levels, releasingenergy (2) lower to higher energy levels, absorbing energy (3)higher to lower energy levels, releasing energy (4) higher to lowerenergy levels, absorbing energy
17. Which Lewis electron-dot structure is drawn correctly forthe atom it represents?
18. When a lithium atom forms a Li+ ion, the lithium atom
(1) gains a proton (3) loses an electron (2) loses a proton (4)gains an electron
19. What is the total number of electrons in the valence shellof an atom of aluminum in the
ground state? (1) 8 (2) 2 (3) 3 (4) 10
20. An electron in an atom moves from the ground state to anexcited state when the energy
of the electron (1) increases (2) decreases (3) remains thesame
21. During a flame test, ions of a specific metal are heated inthe flame of a gas burner. A characteristic color of light isemitted by these ions in the flame when the electrons
(1) emit energy as they move to higher energy levels (2) emitenergy as they return to lower energy levels (3) gain energy asthey move to higher energy levels (4) gain energy as they return tolower energy levels
22. What is the total number of electrons in a Cu+ ion?
(1) 36 (2) 29 (3) 30 (4) 28
Base your answers to questions 23 and 24 on the information andthe bright-line spectra represented below.
Many advertising signs depend on the production of lightemissions from gas-filled glass tubes that are subjected to ahigh-voltage source. When light emissions are passed through aspectroscope, bright-line spectra are produced.
23. Identify the two gases in the unknown mixture. 24. Explainthe production of an emission spectrum in terms of the energystates of an electron.
Atomic Concepts Review Cut from Jan 2007 Jan 2008 Exams
1. Which subatomic particles are located in the nucleus of aneon atom?
(1) electrons and positrons (2) electrons and neutrons q(3)protons and neutrons (4) protons and electrons
2. The total mass of the protons in an atom of gold-198 isapproximately
(1) 79 atomic mass units (2) 119 atomic mass units (3) 198atomic mass units (4) 277 atomic mass units
3. In a calcium atom in the ground state, the electrons thatpossess the least amount of energy are located in the
(1) first electron shell q(2) second electron shell q(3) thirdelectron shell (4) fourth electron shell
4. Which group of atomic models is listed in historical orderfrom the earliest to the most recent? (1) hard-sphere model,wave-mechanical model, electron-shell model (2) hard-sphere model,electron-shell model, wave-mechanical model (3) electron-shellmodel, wave-mechanical model, hard-sphere model (4) electron-shellmodel, hard-sphere model, wave-mechanical model 5 Which isotopicnotation represents an atom of carbon-14?
6. Which isotopic notation identifies a metalloid that ismatched with the corresponding number of protons in each of itsatoms?
7. According to the wave-mechanical model of the atom, electronsin an atom
(1) travel in defined circles (2) are most likely found in anexcited state (3) have a positive charge (4) are located inorbitals outside the
8. What is the total charge of the nucleus of a carbon atom?
(1) 6 (3) +6 (2) 0 (4) +12
9. A sample composed only of atoms having the same atomic numberis classified as
(1) a compound (3) an element (2) a solution (4) an isomer
10. Which two particles each have a mass approximately equal toone atomic mass unit?
(1) electron and neutron (2) electron and positron (3) protonand electron (4) proton and neutron
11. Which electron configuration could represent a strontiumatom in an excited state?
(1) 281871 (3) 281881 (2) 281873 (4) 281882
12. What is the total number of neutrons in an atom
of ? (1) 26 (3) 57 (2) 31 (4) 83
13. What is the total number of electrons in a
Mg +2 ion? (1) 10 (3) 14 (2) 12 (4) 24
14. What was concluded about the structure of the atom as theresult of the gold foil experiment? (1) A positively chargednucleus is surrounded by positively charged particles. (2) Apositively charged nucleus is surrounded by mostly empty space. (3)A negatively charged nucleus is surrounded by positively chargedparticles. (4) A negatively charged nucleus is surrounded by mostlyempty space. 15. An atom is electrically neutral because the (1)number of protons equals the number of electrons (2) number ofprotons equals the number of neutrons (3) ratio of the number ofneutrons to the number of electrons is 1:1 (4) ratio of the numberof neutrons to the number of protons is 2:1 16. How do the energyand the most probable location of an electron in the third shell ofan atom compare to the energy and the most probable location of anelectron in the first shell of the same atom? (1) In the thirdshell, an electron has more energy and is closer to the nucleus.(2) In the third shell, an electron has more energy and is fartherfrom the nucleus. (3) In the third shell, an electron has lessenergy and is closer to the nucleus. (4) In the third shell, anelectron has less energy and is farther from the nucleus. 17. Whatis the net charge on an ion that has 9 protons, 11 neutrons, and 10electrons?
(1) 1+ (3) 1 (2) 2+ (4) 2
18. Which value of an element is calculated using both the massand the relative abundance of each of the naturally occurringisotopes of this
element? (1) atomic number (3) half-life (2) atomic mass (4)molar volume
19. Which two notations represent different isotopes of the sameelement?
Base your answers to questions 20 through 22 on the informationbelow.
The accepted values for the atomic mass and percent naturalabundance of each naturally occurring isotope of silicon are givenin the data table below.
20. Determine the total number of neutrons in an atom of Si-29. 21. Show a correct numerical setup for calculating the atomicmass of Si.  22. A scientist calculated the percent naturalabundance of Si-30 in a sample to be 3.29%. Determine the percenterror for this value.  23. Write one electron configuration foran atom of silicon in an excited state.
16See Also(PDF) Chemistry B Piers Packet - Mrs. Crane's Science ClassPacket... · chemistry b- moles packet designate: _____ hr: _____ page 3 calculate the molar messen for the following combined or diatomic - nvexo.com
Base your answers to questions 24 through 26 on the informationbelow.
24. Identify one piece of information shown in theelectron-shell diagrams that is not shown in the Lewis electron-dotdiagrams.  25. Determine the mass number of the magnesium atomrepresented by the electron-shell diagram.  26. Explain whyLewis electron-dot diagrams are generally more suitable thanelectron-shell diagrams for illustrating chemical bonding. 
Topic 3: Periodic Table Outline 1. The placement of an elementon the Periodic Table gives an indication of the chemical andphysical properties of that element. Elements to the left of thestair step line are metals, and therefore are easily oxidized
(lose electrons) in bonding situations, are good electricalconductors, are shiny, malleable, ductile, and have low ionizationenergies and electronegativities.
Elements to the right of the stair step line, but not in Group18 are nonmetals, and therefore react to gain electrons (getreduced), are not conductors, are dull appearing, brittle, havehigh ionization energies and electronegativities.
Some of the elements along the stair step line have propertiesof both metals and non-metals and are known as metalloids orsemi-metals.
Elements in Group 18 are the noble gases and they are chemicallyinert (unreactive) and have extremely high ionization energies.
2. Elements are arranged in order of increasing atomic number(NOT MASS!) 3. The number of protons in an atom (atomic number)identifies the element. The number of protons in an atom onlychanges through nuclear reactions. 4. The atomic mass is the sum ofprotons and neutrons in the nucleus. The mass number given on theperiodic table is a weighted average of the different isotopes
of that element. Electrons do not significantly add to theatomic mass. 5. Isotopes of an element are identified by the sum ofprotons and neutrons. Isotopes of the same element have the samenumber of protons and a different number of
neutrons. Examples of isotopic notation are: 146C, 14 C,carbon-14, C-14 6. Elements can be classified by their propertiesand their location on the Periodic Table as metals, non-metals,metalloids, and noble gases. 7. Elements may be differentiated bytheir physical properties. Ex: Density, conductivity, malleability,hardness, ductility, solubility 8. Elements may be differentiatedby their chemical properties. Chemical properties describe how anelement behaves in a chemical reaction. 9. Elements are arrangedinto periods and groups. 10. Elements of the same period have thesame number of occupied energy levels.
11. Elements of the same group have the same valenceconfiguration and similar chemical properties. Group 1 elementsother than H are alkali metals. Group 2 elements are alkali earthmetals. Group 17 elements are halogens. Alkali metals, alkali earthmetals, and halogens all are highly reactive and do not existas
free elements in nature (they are all found in compounds). Group18 elements are noble or inert gases. These elements have filledvalence levels and
are do not normally react with other substances. 12. Thesuccession of elements within a group demonstrates characteristictrends in properties. As you progress down a group: atomic radiusincreases. electronegativity decreases. first ionization energydecreases. metallic character increases. 13. The succession ofelements within a period demonstrates characteristic trends inproperties. As you progress across a group from left to right:atomic radius decreases. electronegativity increases. firstionization energy increases. metallic character decreases. 14. Someelements may exist in two or more forms in the same phase. Theseforms differ in their molecular or crystal structure, hence theirdifferent properties. These different forms are called allotropes,Ex: Solid carbon exists in three different forms: graphite, diamond(a network solid) and
coal. Ex: the element oxygen can exist in two different forms:O2 gas and ozone (O3 gas)
Periodic Table Practice Questions
1. Elements in the Periodic Table are arranged according totheir (1) atomic number (3) relative activity (2) atomic mass (4)relative size
2. Elements in a given period of the Periodic Table contain thesame number of (1) protons in the nucleus (3) electrons in theoutermost level (2) neutrons in the nucleus (4) occupied principalenergy levels
3. Atoms of metals tend to
(1) lose electrons and form negative ions (2) lose electrons andform positive ions (3) gain electrons and form negative ions (4)gain electrons and form positive ions
4. Which properties are most common in nonmetals?
(1) low ionization energy and low electronegativity (2) lowionization energy and high electronegativity (3) high ionizationenergy and low electronegativity (4) high ionization energy andhigh electronegativity
5. Which two elements have chemical properties that are mostsimilar?
(1) Cl and Ar (3) K and Ca (2) Li and Na (4) C and N
6. Which of the following Period 4 elements has the mostmetallic characteristics?
(1) Ca (2) Ge (3) As (4) Br 7. If M represents an alkali metal,what is the formula for the compound formed by M and
oxygen? (1) MO2 (2) M2O (3) M2O3 (4) M3O2
8. As the elements in Group 15 are considered in order ofincreasing atomic number, which
sequence in properties occurs? (1) nonmetal, metalloid, metal(3) metal, metalloid, nonmetal (2) metalloid, metal, nonmetal (4)metal, nonmetal, metalloid
9. Which group contains a metalloid?
(1) 1 (2) 11 (3) 15 (4) 18 10. As elements of Group 15 of thePeriodic Table are considered in order from top to
bottom, the metallic character of the atoms of each successiveelement generally (1) increases (2) decreases (3) remains thesame
11. Which statement best describes Group 2 elements as they areconsidered in order from top to bottom of the Periodic Table?
(1) The number of principal energy levels increases, and thenumber of valence electrons increases.
(2) The number of principal energy levels increases, and thenumber of valence electrons remains the same.
(3) The number of principal energy levels remains the same, andthe number of valence electrons increases.
(4) The number of principal energy levels remains the same, andthe number of valence electrons decreases.
12. Which Group 15 element exists as a diatomic molecule atSTP?
(1) phosphorous (3) bismuth (2) nitrogen (4) arsenic
13. Which Group 16 element when combined with hydrogen forms acompound that would
exhibit the strongest hydrogen bonding? (1) selenium (3) oxygen(2) tellurium (4) sulfur
14. Which ion has the largest radius?
(1) Na+ (2) Mg2+ (3) K+ (4) Ca2+ 15. What occurs as the atomicnumber of the elements in Period 2 increases?
(1) The nuclear charge of each successive atom decreases, andthe covalent radius decreases.
(2) The nuclear charge of each successive atom decreases, andthe covalent radius increases.
(3) The nuclear charge of each successive atom increases, andthe covalent radius decreases.
(4) The nuclear charge of each successive atom increases, andthe covalent radius increases.
Periodic Table Cut from Jan 2007 Jan 2008 Exams
1. Which element is a solid at STP and a good conductor ofelectricity?
(1) iodine (3) nickel (2) mercury (4) sulfur
2. Which element has both metallic and nonmetallicproperties?
(1) Rb (3) Si (2) Rn (4) Sr
3. The carbon atoms in graphite and the carbon atoms in diamondhave different
(1) atomic numbers (2) atomic masses (3) electronegativities (4)structural arrangements
4. Atoms of which element have the greatest tendency to gainelectrons?
(1) bromine (3) fluorine (2) chlorine (4) iodine
5. Which statement describes a chemical property of the elementmagnesium?
(1) Magnesium is malleable. (2) Magnesium conducts electricity.(3) Magnesium reacts with an acid. (4) Magnesium has a high boilingpoint.
6. Which statement explains why sulfur is classified as a Group16 element?
(1) A sulfur atom has 6 valence electrons. (2) A sulfur atom has16 neutrons. (3) Sulfur is a yellow solid at STP. (4) Sulfur reactswith most metals.
7. How do the atomic radius and metallic properties of sodiumcompare to the atomic radius and metallic properties ofphosphorus?
(1) Sodium has a larger atomic radius and is more metallic.
(2) Sodium has a larger atomic radius and is lessmetallic. (3)Sodium has a smaller atomic radius
and is more metallic. (4) Sodium has a smaller atomic radius
and is less metallic.
8. Which list of elements consists of metalloids, only?
(1) B, Al, Ga (3) O, S, Se (2) C, N, P (4) Si, Ge, As
9. Which general trend is found in Period 2 on the PeriodicTable as the elements are considered in order of increasing atomicnumber?
(1) decreasing atomic mass (2) decreasing electronegativity (3)increasing atomic radius (4) increasing first ionization energy
10. Which two characteristics are associated with metals? (1)low first ionization energy and low electronegativity (2) low firstionization energy and high electronegativity (3) high firstionization energy and low electronegativity (4) high firstionization energy and high electronegativity 11. Which element ismost chemically similar to chlorine?
(1) Ar (3) Fr (2) F (4) S
12. Which grouping of circles, when considered in order from thetop to the bottom, best represents the relative size of the atomsof Li, Na, K, and Rb, respectively?
13. At STP, which element is brittle and not a conductor ofelectricity?
(1) S (3) Na (2) K (4) Ar
14. An atom of argon rarely bonds to an atom of another elementbecause an argon atom has
(1) 8 valence electrons (2) 2 electrons in the first shell (3) 3electron shells (4) 22 neutrons
15. The elements on the Periodic Table are arranged in order ofincreasing
(1) boiling point (3) atomic number (2) electronegativity (4)atomic mass
16. Which element is classified as a nonmetal?
(1) Be (3) Si (2) Al (4) Cl
17. Solid samples of the element phosphorus can be white, black,or red in color. The variations in color are due to different
(1) atomic masses (2) molecular structures (3) ionizationenergies (4) nuclear charges
18. Lithium and potassium have similar chemical propertiesbecause the atoms of both elements have the same
(1) mass number (2) atomic number (3) number of electron shells(4) number of valence electrons
19. At STP, which list of elements contains a solid, a liquid,and a gas?
(1) Hf, Hg, He (3) Ba, Br2, B (2) Cr, Cl2, C (4) Se, Sn, Sr
Base your answers to questions 20 through 22 on the informationbelow. Elements with atomic numbers 112 and 114 have been producedand their IUPAC names are pending approval. However, an elementthat would be put between these two elements on the Periodic Tablehas not yet been produced. If produced, this element will beidentified by the symbol Uut until an IUPAC name is approved. 20.Draw a Lewis electron-dot diagram for an atom of Uut.  21.Determine the charge of an Uut nucleus. Your response must includeboth the numerical value and the sign of the charge.  22.Identify one element that would be chemically similar to Uut.
Base your answers to questions 23 through 26 on the informationbelow, which describes the proposed discovery of element 118.
In 1999, a nuclear chemist and his team announced they haddiscovered a new element by crashing krypton atoms into lead. Thenew element, number 118, was assigned the name ununoctium and thesymbol Uuo. One possible isotope of ununoctium could have beenUuo-291.
However, the discovery of Uuo was not confirmed because otherscientists could not reproduce the
experimental results published by the nuclear chemist and histeam. In 2006, another team of scientists claimed that theyproduced Uuo. This claim has yet to be confirmed.
Adapted from Discover January 2002
23. Based on atomic number, in which group on the Periodic Tablewould element 118 be placed?  24. What would be the total numberof neutrons present in a theoretical atom of Uuo-291? 25. Whatwould be the total number of electrons present in a theoreticalatom of Uuo-291?  26. Explain why being able to reproducescientific results is an important component of scientificresearch. 
Base your answers to questions 27 through 30 on the informationbelow.
The table below lists physical and chemical properties of sixelements at standard pressure that correspond to known elements onthe Periodic Table. The elements are identified by the codeletters, D, E, G, J, L, and Q.
27. What is the total number of elements in the Properties ofSix Elements at Standard Pressure table that are solids at STP? 28. An atom of element G is in the ground state. What is the totalnumber of valence electrons in this atom?  29. Letter Zcorresponds to an element on the Periodic Table other than the sixlisted elements. Elements G, Q, L, and Z are in the same group onthe Periodic Table, as shown in the diagram below.
Based on the trend in the melting points for elements G, Q, andL listed in the Properties of Six Elements at Standard Pressuretable, estimate the melting point of element Z, in degrees Celsius. 30. Identify, by code letter, the element that is a noble gasin the Properties of Six Elements at Standard Pressure table.
Topic 4: Formulas & Names, Equations, Moles, Molar Mass,& Types of Reactions Outline
1. A compound is a substance composed of two or more differentelements that are chemically combined in a fixed proportion. Achemical compound can only be broken down by chemical means. 2.Chemical compounds can be represented by a specific formula andassigned a name based on the IUPAC system. 3. Types of chemicalformulas include empirical, molecular, and structural. Empiricalformulas show elements in their simplest whole number ratios. Thismay or may
not be the same as the molecular formula. Molecular formulasshow the actual number of atoms per element in a single molecule.Structural formulas show the number of each type of atom as well astheir physical
arrangement. 4. All chemical reactions show a conservation ofmass, energy and charge. 5. A balanced chemical equation representsconservation of atoms. 6. The coefficients in a balanced chemicalequation can be used to determine mole ratios in the reaction, andcan further be used to predict relationships about amounts betweenproducts and reactants. 7. The molar mass of a substance is the sumof the atomic masses of its atoms. The molar mass (gram formulamass) equals the mass of one mole of that substance. 8. The percentcomposition by mass of each element in a compound can be calculatedmathematically. 9. Types of chemical reactions include synthesis,decomposition, single replacement, and double replacement.
Equations & Stoichiometry Practice Questions
1. Which substance has the greatest molecular mass? (1) H2O2 (2)NO (3) CF4 (4) I2
2. What is the gram formula mass of Ca(OH)2? (1) 29 g (2) 34 g(3) 57 g (4) 74 g
3. What is the total number of moles of atoms present in 1 gramformula mass of
Pb(C2H3O2)2? (1) 9 (2) 14 (3) 3 (4) 15
4. The percent by mass of carbon in HC2H3O2 is equal to
5. What is the empirical formula of C3H6?
(1) CH (2) CH2 (3) CH3 (4) CH6 6. The name of the compound KClO2is potassium
(1) hypochlorite (3) chlorate (2) chlorite (4) perchlorate
7. Which formula is correct for ammonium sulfate?
(1) NH4SO4 (2) (NH4)2SO4 (3) NH4(SO4)2 (4) (NH)3(SO4)2 8. Themolecular formula of a compound is represented by X3Y6. What is theempirical
formula of this compound? (1) X3Y (2) X2Y (3) XY2 (4) XY
9. The number of moles of molecules in a 12.0-gram samples ofCl2 is
mole (2) 0.710.12
mole (3) 12.0 moles (4) 12.0 x 35.5 moles
10. What is the total number of moles of sulfur atoms in 1 moleof Fe2(SO4)3?
(1) 1 (2) 15 (3) 3 (4) 17 11. Given the unbalanced equation:
___ CaSO4 + ___ AlCl3 ___ Al2(SO4)3 + ___ CaCl2
What is the coefficient of Al2(SO4)3 when the equation iscompletely balanced using the smallest whole-numbercoefficients?
(1) 1 (2) 2 (3) 3 (4) 4
12. Given the unbalanced equation:
___ Al (s) + ___ O2 (g) ___ Al2O3 (s)
When this equation is correctly balanced using smallest wholenumbers, what is the coefficient of O2 (g)?
(1) 6 (2) 2 (3) 3 (4) 4 13. Given the reaction:
4 NH3 + 5 O2 4 NO + 6 H2O
What is the total number of moles of NO produced when 1.0 moleof O2 is completely consumed? (1) 1.0 mole (2) 1.2 moles (3) 0.80mole (4) 4.0 moles
14. Given the equation:
H2 (g) + Cl2 (g) 2 HCl (g)
What is the total number of moles of HCl (g) produced when 3moles of H2 (g) is completely consumed? (1) 5 moles (2) 2 moles (3)3 moles (4) 6 moles
Formulas, Equations & Stoichiometry Review Cut from Jan 2007Jan 2008 Exams
1. Which equation shows conservation of atoms?
(1) H2 + O2 H2O (2) H2 + O2 2 H2O (3) 2 H2 + O2 2 H2O (4) 2 H2 +2 O2 2 H2O
2. Which substance can be broken down by a chemical change?
(1) antimony (3) hexane (2) carbon (4) sulfur
3. What is the gram formula mass of Ca3(PO4)2?
(1) 248 g/mol (3) 279 g/mol (2) 263 g/mol (4) 310 g/mol
4. In which compound is the ratio of metal ions to nonmetal ions1 to 2?
(1) calcium bromide (2) calcium oxide (3) calcium phosphide (4)calcium sulfide
5. Given the balanced equation representing a
reaction: 2CO(g) + O2(g) 2CO2(g)
What is the mole ratio of CO(g) to CO2(g) in this reaction? (1)1:1 (3) 2:1 (2) 1:2 (4) 3:2
6. Given the balanced equation representing a
reaction: H+(aq) + OH(aq) H2O(l) + 55.8 kJ
In this reaction there is conservation of (1) mass, only (2)mass and charge, only (3) mass and energy, only (4) mass, charge,and energy
7. Which polyatomic ion contains the greatest number of oxygenatoms?
(1) acetone (3) hydroxide (2) carbonate (4) peroxide
8. Which formula represents an ionic compound?
(1) H2 (3) CH3OH (2) CH4 (4) NH4Cl
9. What is the total number of different elements present inNH4NO3?
(1) 7 (3) 3 (2) 9 (4) 4
10. Which formula represents lead (II) chromate?
(1) PbCrO4 (3) Pb2CrO4 (2) Pb(CrO4)2 (4) Pb2(CrO4)3
11. Which particle diagram represents a sample of one compound,only?
12. An atom in the ground state contains a total of 5 electrons,5 protons, and 5 neutrons. Which Lewis electron-dot diagramrepresents this atom?
13. Given the balanced equation representing
the reaction between propane and oxygen: C3H8 + 5O2 3CO2 +4H2O
According to this equation, which ratio of oxygen to propane iscorrect?
14. Which substance can be decomposed by chemical means?
(1) tungsten (3) krypton (2) antimony (4) methane
15. Given the balanced equation representing a reaction:
4NH3 + 5O2 4NO + 6H2O What is the minimum number of moles of O2that are needed to completely react with 16 moles of NH3? (1) 16mol (3) 64 mol (2) 20. mol (4) 80. mol
16. Element X reacts with iron to form two
different compounds with the formulas FeX and Fe2X3. To whichgroup on the Periodic Table does element X belong? (1) Group 8 (3)Group 13 (2) Group 2 (4) Group 16
17. The molar mass of Ba(OH)2 is
(1) 154.3 g (3) 171.3 g (2) 155.3 g (4) 308.6 g
18. Given the balanced equation representing a
reaction: H2SO4(aq) + 2KOH(aq) K2SO4(aq) + 2H2O(l)
Which type of reaction is represented by this equation? (1)decomposition (2) neutralization (3) single replacement (4)synthesis
19. A hydrated compound contains water molecules within itscrystal structure. The percent composition
by mass of water in the hydrated compound CaSO42H2O has anaccepted value of 20.9%. A student did an experiment and determinedthat the percent composition by mass of water in CaSO42H2O was21.4%. Calculate the percent error of the students experimentalresult. Your response must include both a correct numerical setupand the calculated result. 
20. Write the empirical formula for the compound C8H18. 
Base your answers to questions 21 through 23 on the informationbelow.
Some dry chemicals can be used to put out forest fires. One ofthese chemicals is NaHCO3. When NaHCO3(s) is heated, one of theproducts is CO2(g), as shown in the balanced equation below.
2 NaHCO3(s) + heat Na2CO3(s) + H2O(g) + CO2(g)
21. Show a correct numerical setup for calculating the percentcomposition by mass of carbon in the product Na2CO3. 
22. Identify whether the reaction is endothermic or exothermic. 23. Determine the total number of moles of CO2(g) produced when7.0 moles of NaHCO3(s) is
completely reacted. 
moles 24. Balance this chemical equation:  _____ S(s) + _____KClO3(s) _____SO2(g) + _____ KCl(s) + energy Base your answers toquestions 25 through 27 on the information below.
Rust on an automobile door contains Fe2O3(s). The balancedequation representing one of the reactions between iron in the doorof the automobile and oxygen in the atmosphere is given below.
4Fe(s) + 3O2(g) 2Fe2O3(s)
25. Identify the type of chemical reaction represented by thisequation.  26. Determine the gram-formula mass of the product ofthis reaction.  27. Write the IUPAC name for Fe2O3. 
Base your answers to questions 28 through 30 on the informationbelow.
Ozone gas, O3, can be used to kill adult insects in storage binsfor grain without damaging the grain. The ozone is produced fromoxygen gas, O2, in portable ozone generators located near thestorage bins. The concentrations of ozone used are so low that theydo not cause any environmental damage. This use of ozone is saferand more environmentally friendly than a method that usedbromomethane, CH3Br. However, bromomethane was more effective thanozone because CH3Br killed immature insects as well as adultinsects.
Adapted From: The Sunday Gazette (Schenectady, NY) 3/9/03 28.Determine the total number of moles of CH3Br in 19 grams of CH3Br(gram-formula mass =
95 grams/mol). 
29. Given the balanced equation for producing bromomethane:
Br2 + CH4 CH3Br + HBr
Identify the type of organic reaction shown. 
30. Based on the information in the passage, state one advantageof using ozone instead of bromomethane for insect control in grainstorage bins. 
Base your answers to questions 31 through 33 on the informationbelow.
A hydrate is a compound that has water molecules within itscrystal structure. The formula for the hydrate CuSO45H2O(s) showsthat there are five moles of water for every one mole of CuSO4(s).When CuSO45H2O(s) is heated, the water within the crystals isreleased, as represented by the balanced equation below.
CuSO45H2O(s) CuSO4(s) + 5H2O(g)
A student first masses an empty crucible (a heat-resistantcontainer). The student then
masses the crucible containing a sample of CuSO45H2O(s). Thestudent repeatedly heats and masses the crucible and its contentsuntil the mass is constant. The students recorded experimental dataand calculations are shown below.
31. Identify the total number of significant figures recorded inthe calculated mass of CuSO45H2O(s).  32. In the space below,use the students data to show a correct numerical setup forcalculating the percent
composition by mass of water in the hydrate.  33. Explain whythe sample in the crucible must be heated until the constant massis reached. 
Topic 5: Bonding Outline 1. Chemical compounds are formed whenatoms are bonded together. Breaking a chemical bond is anendothermic process. Forming a chemical bond is an exothermicprocess. Compounds have less potential energy than the individualatoms they are formed from. 2. Two major categories of compoundsare ionic and molecular (covalent) compounds. Ionic compounds tendto be a metal bonding with a nonmetal; or a metal with apolyatomic
ion Molecular (covalent) compounds tend to be two or morenonmetals combined. 3. Compounds can be differentiated by theirchemical and physical properties. Ionic substances have highmelting and boiling points, form crystals, dissolve in water
(dissociate), and conduct electricity in solution and asliquids. Covalent or molecular substances have lower melting andboiling points, do not conduct
electricity. 4. Atoms gain a stable electron configuration bybonding with other atoms. Atoms are stable when they have a fullvalence level. Most atoms need 8 electrons to fill their valencelevel. H and He only need 2 electrons to fill their valence level.The noble gases (group 18) have filled valence levels. They do notnormally bond with other
atoms. 5. Chemical bonds are formed when valence electrons are:Transferred from one atom to another ionic. Shared between atomscovalent. Mobile in a free moving sea of electrons metallic. 6. Inmultiple (double or triple) covalent bonds more than 1 pair ofelectrons are shared between two atoms. oxygen and its family(group 16) form double bonds with each other (O2) nitrogen and itsfamily (group 15) form triple bonds with each other (NH3) carboncan form double and triple bonds with itself & group 16 and 15elements (ex: CO2)
7. Polarity of a molecule can be determined by its shape and thedistribution of the charge. Polar molecules have an asymmetrical(uneven) distribution of electrons in them. As a result, polarmolecules have (+) and (-) charged ends. Water is the most commonsubstance composed of polar molecules; O end is (-), H ends are
(+). Nonpolar molecules have symmetrical (even) distribution ofelectrons in them. Polar substances are dissolved only by anotherpolar substance. Non-polar substances are
dissolved only by other non-polar substances. 8. Theelectronegativity difference between two bonded atoms can determinethe type of bond and its polarity. 0.0 = non-polar covalent 0.0-1.7 = polar covalent 1.7+ = ionic 9. Bonding guidelines: Metalsreact with nonmetals to form ionic compounds. Nonmetals bond withnonmetals to form covalent compounds (molecules). Ionic compoundswith polyatomic ions have both ionic and covalent bonds. 10.Intermolecular forces allow different particles to be attracted toeach other to form solids and liquids. Hydrogen bonds are anexample of a strong IMF between polar molecules. Hydrogen bondsexist between atoms of hydrogen on one molecule and atoms ofeither
oxygen, fluorine, or nitrogen on a neighboring molecule.Substances with hydrogen bonds tend to have much higher melting andboiling points than
those without hydrogen bonds. Water is one such substanceOrdinary polar molecules simply attract each other as theiroppositely charged ends line up. Nonpolar molecules use weakVanderWaals forces of attraction and as a result tend to have
lower melting points, and higher vapor pressures. 11. Metallicbonding occurs between atoms of metal. The valence electrons areloosely held by all atoms in a mobile sea of valence electrons.This type of bonding accounts for some of the unique properties ofmetals, such as their
ability to conduct electricity, luster, and malleability. 12.Physical properties of a substance can be explained in terms ofchemical bonds and intermolecular forces. These includeconductivity, malleability, solubility, ductility, hardness,melting point and boiling point, vapor pressure.
Bonding Practice Questions
1. The forces between atoms that create chemical bonds are theresult of interactions between (1) nuclei (3) protons and electrons(2) electrons (4) protons and nuclei
2. According to Reference Table S, which sequence correctlyplaces the elements in order of increasing ionization energy? (1) HLi Na K (3) O S Se Te (2) I Br Cl F (4) H Be Al Ga
3. Electronegativity is a measure of an atoms ability to
(1) attract the electrons in the bond between the atom andanother atom (2) repel the electrons in the bond between the atomand another atom (3) attract the protons of another atom (4) repelthe protons of another atom
4. If the electronegativity difference between the elements incompound NaX is 2.0, what is
element X? (1) bromine (2) chlorine (3) fluorine (4) oxygen
5. An element with an electronegativity of 0.9 bonds with anelement with an
electronegativity of 3.1. Which phrase best describes the bondbetween these elements? (1) mostly ionic in character and formedbetween two nonmetals (2) mostly ionic in character and formedbetween a metal and a nonmetal (3) mostly covalent in character andformed between two nonmetals (4) mostly covalent in character andformed between a metal and a nonmetal
6. Which type of bond exists between an atom of carbon and anatom of fluorine?
(1) ionic (2) metallic (3) polar covalent (4) nonpolar covalent7. Which pair of atoms is held together by a covalent bond?
(1) HCl (2) LiCl (3) NaCl (4) KCl 8. Which substance containsnonpolar covalent bonds?
(1) H2 (2) H2O (3) Ca(OH)2 (4) CaO 9. Given the reaction: Cl (g)+ Cl (g) Cl2 (g) + energy
Which statement best describes the reaction? (1) A bond isformed and energy is absorbed. (2) A bond is formed and energy isreleased. (3) A bond is broken and energy is absorbed. (4) A bondis broken and energy is released.
10. The primary forces of attraction between water molecules inH2O (l) are
(1) ionic bonds (3) molecule-ion attractions (2) hydrogen bonds(4) van der Waals forces
11. Which structure represents a polar molecule?
(1) H H (3)
12. Which electron dot diagram represents a molecule that has apolar covalent bond? 1) 2)
Bonding Review Cut from Jan 2007 Jan 2008 Exams
1. Given the balanced equation:
I + I I2 Which statement describes the process represented bythis equation? (1) A bond is formed as energy is absorbed. (2) Abond is formed and energy is released. (3) A bond is broken asenergy is absorbed. (4) A bond is broken and energy is released. 2.An oxygen molecule contains a double bond because the two atoms ofoxygen share a total of
(1) 1 electron (3) 3 electrons (2) 2 electrons (4) 4electrons
3. A double carbon-carbon bond is found in a molecule of
(1) pentane (3) pentyne (2) pentene (4) pentanol
4. At STP, fluorine is a gas and bromine is a liquid because,compared to fluorine, bromine has
(1) stronger covalent bonds (2) stronger intermolecular forces(3) weaker covalent bonds (4) weaker intermolecular forces
5. Which term indicates how strongly an atom attracts theelectrons in a chemical bond?
(1) alkalinity (2) atomic mass (3) electronegativity (4)activation energy
6. Magnesium nitrate contains chemical bonds that are
(1) covalent, only (2) ionic, only (3) both covalent and ionic(4) neither covalent nor ionic
7. A solid substance is an excellent conductor of electricity.The chemical bonds in this substance are most likely (1) ionic,because the valence electrons are shared between atoms (2) ionic,because the valence electrons are mobile (3) metallic, because thevalence electrons are stationary (4) metallic, because the valenceelectrons are mobile
8. When sodium and fluorine combine to produce the compound NaF,the ions formed have the same electron configuration as atomsof
(1) argon, only (2) neon, only (3) both argon and neon (4)neither argon nor neon
9. Atoms of which element have the greatest tendency to gainelectrons?
(1) bromine (3) fluorine (2) chlorine (4) iodine
10. Which polyatomic ion contains the greatest number of oxygenatoms?
(1) acetate (3) hydroxide (2) carbonate (4) peroxide
11. Which formula represents an ionic compound? (1) H2 (3) CH3OH(2) CH4 (4) NH4Cl
12. Which liquid has the highest vapor pressure at 75C?
(1) ethanoic acid (3) propanone (2) ethanol (4) water
13. Given the balanced equation representing a reaction:
Cl2(g) Cl(g) + Cl(g) What occurs during this change? (1) Energyis absorbed and a bond is broken. (2) Energy is absorbed and a bondis formed. (3) Energy is released and a bond is broken. (4) Energyis released and a bond is formed. 14. At standard pressure, acertain compound has a low boiling point and is insoluble in water.At STP, this compound most likely exists as
(1) ionic crystals (2) metallic crystals (3) nonpolar molecules(4) polar molecules
15. Which group on the Periodic Table of the Elements containselements that react with oxygen to form compounds with the generalformula X2O?
(1) Group 1 (3) Group 14 (2) Group 2 (4) Group 18
16. Which two substances are covalent compounds?
(1) C6H12O6(s) and KI(s) (2) C6H12O6(s) and HCl(g) (3) KI(s) andNaCl(s) (4) NaCl(s) and HCl(g)
17. Which compound has hydrogen bonding between itsmolecules?
(1) CH4 (3) KH (2) CaH2 (4) NH3
18. Which Lewis electron-dot diagram correctly represents ahydroxide ion?
19. Explain, in terms of electronegativity, why a PCl bond in amolecule of PCl5 is more polar than a PS bond in a molecule ofP2S5.  Base your answers to questions 20 and 21 on theinformation below.
The graph below shows the relationship between boiling point andmolar mass at standard pressure for pentane, hexane, heptane, andnonane.
20. Octane has a molar mass of 114 grams per mole. According tothis graph, what is the boiling point of octane at standardpressure?  ________________________ 21. State the relationshipbetween molar mass and the strength of intermolecular forces forthe selected alkanes. 
Base your answers to questions 22 through 24 on the informationbelow.
The particle diagrams below represent the reaction between twononmetals, A2 and Q2.
22. Using the symbols A and Q, write the chemical formula of theproduct.  ________________ 23. Identify the type of chemicalbond between an atom of element A and an atom of element Q.  24.Compare the total mass of the reactants to the total mass of theproduct.  25. Explain, in terms of molecular structure ordistribution of charge, why a molecule of methane is nonpolar. 26. A liquid boils when the vapor pressure of the liquid equals theatmospheric pressure on the surface of the liquid. Using ReferenceTable H, determine the boiling point of water when the atmosphericpressure is 90. kPa. 
Base your answers to questions 27 through 30 on the informationbelow. Have you ever seen an insect called a water strider skatingacross the surface of a calm pond? Have you ever floated a sewingneedle on the water in a glass? If you have, then youve observedone of waters many amazing properties. Waters surface tension keepsthe water strider and the sewing needle from sinking into thewater. Simply stated, the surface tension is due to the forces thathold the water molecules together. Without these intermolecularforces, the water strider and the sewing needle would sink belowthe surface of the water. The surface tension of water at varioustemperatures is given in the data table below.
27. On a piece of graph paper, plot the data from the datatable. Circle and connect the five points.  28. According toyour graph, what is the surface tension of water at 60.C? ___________ mN/m 29. State the relationship between the surfacetension and the temperature of water.  30. The surface tensionof liquid tetrachloromethane, CCl4, at 25C is 26.3 millinewtons/meter (mN/m). Compare the intermolecular forces between moleculesof CCl4 to the intermolecular forces between molecules of water,H2O, at 25C. 
Topic 6 Overview Topic 6A: Heat & Temperature 1. Energy canexist in different forms chemical, electrical, electromagnetic,thermal, mechanical, nuclear. Stored energy is referred to aspotential energy. Energy of motion is kinetic energy. 2. The Law ofConservation of Energy states that energy can not be lost ordestroyed, only changed from one form to another. 3. Heat is atransfer of energy (often but not always thermal energy) from abody of higher temperature to a body of lower temperature. 4.Temperature is a measure of the average kinetic energy of theparticles in a sample. Temperature is NOT a form of energy andshould not be confused with heat. 5. The concepts of kinetic andpotential energy can be used to explain physical processes such asfusion (melting), solidification (freezing), vaporization (boiling,evaporation), condensation, sublimation, and deposition. 6.Processes that are exothermic give off heat energy. This typicallycauses the surrounding environment to become warmer. 7. Processesthat are endothermic absorb energy. This typically causes thesurrounding environment to become colder.
Topic 6A: Heat & Temperature Outline 1. Temperature is ameasure of the average kinetic energy of the particles in a sampleof matter. Kinetic energy is energy due to motion. So astemperature increases, the particles move
faster, on average. Temperature does NOT depend on the mass ofthe sample.
2. Temperature scales used by chemists are the Celsius andKelvin scales. The freezing point of water is a reference pointoften used in science, and is referred to
as standard temperature. Its value is 0oC or 273 K, and is notedon Table A. The boiling point of water is 100oC or 373 K.Converting from Co to K: K = Co + 273 (on Table T)
3. Heat is a form of energy and IS NOT the same as temperature.Heat is dependent on mass. There is more heat in an iceberg that isat 00C than a cup
full of boiling water. Heat can be transferred from onesubstance to another when their particles are in
contact (when the objects touch). Heat will move from the objectwith more particle KE (higher temp) to the one with less.
The amount of heat needed to cause a temperature change isdependent on the mass of the sample, its specific heat and theamount of temperature change: q = m c T (Table T) When heat isabsorbed to cause a temperature change, it is resulting in a changein KE of particles.
The amount of heat needed to cause a phase change can becalculated using the q = mHf (melting), or q = mHv (boiling) (TableT). When heat is added to cause a phase change, it is causing achange in intermolecular forces between particles.
The values for water are on Table B. 4. Heat of fusion (Hf) isthe energy needed to convert one gram of a substance from solid toliquid. 5. Heat of vaporization (Hv) is the energy needed toconvert one gram of a substance from liquid to gas. 6. Specificheat (C) is the energy required to raise one gram of a substance 1degree (Celcius or Kelvin). The specific heat of liquid water is 1cal/g*J or 4.2 J/g*K.
7. The three phases of matter are solid, liquid and gas. Eachhas its own properties. Solids have a constant volume and shape.Particles are held in a rigid, crystalline structure. Liquids havea constant volume but a changing shape. Particles are mobile butstill held
together by strong attraction. Gases have no set volume orshape. They will completely fill any closed contained.Particles
have largely broken free of the forces holding them together.The phase a substance is in is dependent on the temperature.Melting points and boiling
points are on Table S (in Kelvin degrees). 8. Phase changes area type of physical change. If they are changes that involve heatbeing absorbed, they are endothermic changes. Endothermic phasechanges are melting, boiling, evaporating and subliming (sl).Opposite type of phase changes (freezing, condensing, depositing)are exothermic.
9. A heating curve (or cooling curve) traces the changes intemperature of a substance as it changes from solid to liquid togas (or gas to liquid to solid). When the substance undergoes aphase change, there is no change in temperature. The line
flattens until the phase change is complete. When a phase changeis occurring, the potential energy of the substance changeswhile
kinetic energy remains the same. As temperature increases,kinetic energy increases. 10. The amount of heat involved in somechemical changes is shown on Table I, called heat of reaction or H.If the value is negative, the reaction is exothermic. This can beexpressed as a potential energy diagram. If the energy is writteninto the equation, and is on the reactants side, the reactionis
endothermic. H is the difference between the energy stored inthe products (PE) and the potential
energy of the reactants. 11. Breaking bonds is ALWAYSendothermic, and forming bonds is ALWAYS exothermic. I + I I2 Bondis forming, I atoms are become stable by bonding, so theyrelease
energy (Exo) H2 H + H Bond is breaking, requires energy in orderto put atoms in unbonded
Heat and Temperature Cut from Jan 2007 Jan 2008 Exams
1. Given the balanced equation: I + I I2
Which statement describes the process represented by thisequation? (1) A bond is formed as energy is absorbed. (2) A bond isformed and energy is released. (3) A bond is broken as energy isabsorbed. (4) A bond is broken and energy is released. 2. Whichterm is defined as a measure of the average kinetic energy of theparticles in a sample?
(1) temperature (3) thermal energy (2) pressure (4) chemicalenergy
3. Which term refers to the difference between the potentialenergy of the products and the potential energy of the reactantsfor any chemical change? (1) heat of deposition (2) heat of fusion(3) heat of reaction (4) heat of vaporization 4. Which kelvintemperature is equal to 56C?
(1) 329 K (3) 217 K (2) 217 K (4) 329 K
5. Which reaction releases the greatest amount of energy per 2moles of product? (1) 2CO(g) + O2(g) 2CO2(g) (2) 4Al(s) + 3O2(g)2Al2O3(s) (3) 2H2(g) + O2(g) 2H2O(g) (4) N2(g) + 3H2(g) 2NH3(g)
Use the reaction shown below to answer questions 6 and 7.
6. Draw a potential energy diagram for this reaction.  7.Determine the total amount of energy released when 2.50 moles ofpropane is completely reacted with oxygen. 
8. Given the balanced equation representing a reaction: N2(g) +O2(g) + 182.6 kJ 2NO(g) Draw a potential energy diagram for thisreaction.  Base your answers to questions 9 through 11 on theinformation below.
A 5.00-gram sample of liquid ammonia is originally at 210. K.The diagram of the partial heating curve below represents thevaporization of the sample of ammonia at standard pressure due tothe addition of heat. The heat is not added at a constant rate.
Some physical constants for ammonia are shown in the data tablebelow.
9. Calculate the total heat absorbed by the 5.00-gram sample ofammonia during time interval AB. Your response must include both acorrect numerical setup and the calculated result.  10. Describewhat is happening to both the potential energy and the averagekinetic energy of the molecules in the ammonia sample during timeinterval BC. Your response must include both potential energy andaverage kinetic energy.  11. Determine the total amount of heatrequired to vaporize this 5.00-gram sample of ammonia at itsboiling point. 
Base your answers to questions 12 through 14 on the informationbelow.
A 100.0-gram sample of NaCl(s) has an initial temperature of 0C.A chemist measures the temperature of the sample as it is heated.Heat is not added at a constant rate. The heating curve for thesample is shown below.
12. Determine the temperature range over which the entire NaClsample is a liquid.  13. Identify one line segment on the curvewhere the average kinetic energy of the particles of the NaClsample is changing.  14. Identify one line segment on the curvewhere the NaCl sample is in a single phase and capable ofconducting electricity. 
Base your answers to questions 15 and 16 on the informationbelow. A student performed an experiment to determine the totalamount of energy stored in a peanut. The
accepted value for the energy content of a peanut is 30.2kilojoules per gram. The student measured 100.0 grams of water intoa metal can and placed the can on a ring stand, as shown in thediagram below. The peanut was attached to a wire suspended underthe can. The initial temperature of the water was recorded as22.0C. The peanut was ignited and allowed to burn. When the peanutfinished burning, the final water temperature was recorded as57.0C. The students experimental value for the energy content ofthis peanut was 25.9 kilojoules per gram.
15. Calculate the total amount of heat absorbed by the water.Your response must include both a correct numerical setup and thecalculated result.  16. Determine the students percent error forthe energy content of this peanut. 
Base your answers to questions 17 through 20 on the informationbelow.
The temperature of a sample of a substance is increased from20.C to 160.C as the sample absorbs heat at a constant rate of 15kilojoules per minute at standard pressure. The graph belowrepresents the relationship between temperature and time as thesample is heated.
17. What is the boiling point of this sample?  18. Draw atleast nine particles in the box, showing the correct particlearrangement of this sample during the first minute of heating. 19. What is the total time this sample is in the liquid phase,only?  20. Determine the total amount of heat required tocompletely melt this sample at its melting point. 
Topic 6B: Reaction Rate & Equilibrium Outline 1. Collisiontheory states that a reaction is most likely to occur if reactantparticles collide with the proper energy and orientation. This issometimes called an effective collision.
2. The rate of a chemical reaction depends on temperature,concentration, nature of the reactants, surface area and thepresence of a catalyst.
3. Energy absorbed or released by a chemical reaction can berepresented by a potential energy diagram.
4. The amount of energy released or absorbed during a chemicalreaction is the heat of reaction. Heat of reaction equals the PE ofthe products PE of reactants. Positive heat of reaction implies anendothermic reaction. Negative heat of reaction implies anexothermic reaction.
5. A catalyst provides an alternative pathway for a chemicalreaction. The catalyst lowers reaction the activation energyrequired to start up the reaction. Adding a catalyst increases therate of the forward and reverse reactions equally, so there is noshift
in equilibrium. Know how the use of a catalyst affects the PEdiagram.
6. Entropy is a measure of the randomness or disorder in asystem. A system with greater disorder has greater entropy.
7. Systems in nature tend to undergo changes towards lowerenergy (tend to be exothermic) and higher entropy.
8. At equilibrium the rate of the forward reaction equals therate of the reverse reaction. This state can only be achieved IFthe system (container) is closed and the conditions of Temp
and Pressure are held steady.
9. The measurable quantities of reactants and products remainconstant at equilibrium. This does NOT mean the amounts of productsand reactants is the same as each other, but
rather that the amounts are no longer changing.
10. Types of equilibrium include chemical, phase and solution.Solutions that are saturated represent an equilibrium between theprocesses of dissolving and
precipitating. An example of a phase equilibrium would be thesimultaneous melting and freezing of water if
the system is held at 0oC.
11. LeChateliers principle can be used to predict the effect ofstress on a system in equilibrium. Stresses include a change inpressure, volume, concentration, and temperature. You should beable to predict if a shift left or a shift right occurs due to aparticular stress.
Rate of Reaction & Equilibrium Cut from Jan 2007 Jan 2008Exams
1. Given the equation representing a phase change atequilibrium:
Which statement is true? (1) The forward process proceeds fasterthan the reverse process. (2) The reverse process proceeds fasterthan the forward process. (3) The forward and reverse processesproceed at the same rate. (4) The forward and reverse processesboth stop. 2. A 5.0-gram sample of zinc and a 50.-milliliter sampleof hydrochloric acid are used in a chemical reaction. Whichcombination of these samples has the fastest reaction rate? (1) azinc strip and 1.0 M HCl(aq) (2) a zinc strip and 3.0 M HCl(aq) (3)zinc powder and 1.0 M HCl(aq) (4) zinc powder and 3.0 M HCl(aq) 3.For a given reaction, adding a catalyst increases the rate of thereaction by (1) providing an alternate reaction pathway that has ahigher activation energy (2) providing an alternate reactionpathway that has a lower activation energy (3) using the samereaction pathway and increasing the activation energy (4) using thesame reaction pathway and decreasing the activation energy 4. Giventhe equation representing a reaction at equilibrium:
Which change causes the equilibrium to shift to the right? (1)decreasing the concentration of H2(g) (2) decreasing the pressure(3) increasing the concentration of N2(g) (4) increasing thetemperature
5. Given the equation representing a system at equilibrium:
At which temperature does this equilibrium exist at 101.3kilopascals?
(1) 0 K (3) 32 K (2) 0C (4) 273C
6. Which statement must be true when solution equilibriumoccurs? (1) The solution is at STP. (2) The solution issupersaturated. (3) The concentration of the solution remainsconstant. (4) The masses of the dissolved solute and theundissolved solute are equal. 7. Which statement must be true forany chemical reaction at equilibrium? (1) The concentration of theproducts is greater than the concentration of the reactants. (2)The concentration of the products is less than the concentration ofthe reactants. (3) The concentration of the products and theconcentration of the reactants are equal. (4) The concentration ofthe products and the concentration of the reactants are constant.8. Given the balanced equation representing a reaction at 101.3 kPaand 298 K:
Which statement is true about this reaction? (1) It isexothermic and H equals 91.8 kJ. (2) It is exothermic and H equals+91.8 kJ. (3) It is endothermic and H equals 91.8 kJ. (4) It isendothermic and H equals +91.8 kJ. 9. Which balanced equationrepresents a phase
10. Given the system at equilibrium:
Which changes occur when O2(g) is added to this system? (1) Theequilibrium shifts to the right and the concentration of PCl3(g)increases. (2) The equilibrium shifts to the right and theconcentration of PCl3(g) decreases. (3) The equilibrium shifts tothe left and the concentration of PCl3(g) increases. (4) Theequilibrium shifts to the left and the concentration of PCl3(g)decreases.
11. In terms of energy and entropy, systems in nature tend toundergo changes toward (1) higher energy and higher entropy (2)higher energy and lower entropy (3) lower energy and higher entropy(4) lower energy and lower entropy
12. Explain, in terms of collision theory, why the rate of achemical reaction increases with an increase in temperature. Base your answers to questions 13 through 15 on the informationbelow.
A beaker contains 100.0 milliliters of a dilute aqueous solutionof ethanoic acid at equilibrium. The equation below represents thissystem.
13. Compare the rate of the forward reaction to the rate of thereverse reaction for this system.  14. Describe what happens tothe concentration of H+(aq) when 10 drops of concentratedHC2H3O2(aq) are added to this system.  15. Draw a structuralformula for ethanoic acid. 
Base your answer to question 16 on the information below.
Hand Blasters is a toy that consists of a set of two ceramicballs, each coated with a mixture of sulfur and potassium chlorate,KClO3. When the two balls are struck together, a loud popping noiseis produced as sulfur and potassium chlorate react with each other.16. Identify one source of the activation energy for this reaction. Base your answers to questions 17 through 18 on the reactionrepresented by the balanced equation below.
2H2(g) + O2(g) 2H2O(l) + 571.6 kJ
17. Identify the information in this equation that indicates thereaction is exothermic.  18. Explain why the entropy of thesystem decreases as the reaction proceeds.  Base your answers toquestion 19 on the information below.
The equilibrium equation below is related to the manufacture ofa bleaching solution. In this equation, Cl(aq) means that chlorideions are surrounded by water molecules.
19. Explain, in terms of collision theory, why increasing theconcentration of Cl2(g) increases the concentration of OCl(aq) inthis equilibrium system. 
Base your answers to questions 20 and 21 on the informationbelow.
A gasoline engine burns gasoline in the presence of excessoxygen to form carbon dioxide and water. The main components ofgasoline are isomers of octane. A structural formula of octane isshown below.
One isomer of octane is 2,2,4-trimethylpentane.
20. In the space in your answer booklet, draw a structuralformula for 2,2,4-trimethylpentane.  21. Explain, in terms ofthe arrangement of particles, why the entropy of gasoline vapor isgreater than the entropy of liquid gasoline. 
Topic 7: Water & Solutions 1. Water has some unusualproperties. The bonds between H and O inside a water molecule arepolar covalent. Due to its structure, it is a polar molecule. Thismeans it has an uneven distribution of
electrons in it. The O end is (-) and the H ends are (+). Drawits Lew is dot structure here! ______________ Water is actually aVERY polar substance. As a result, it uses the very strong typeof
intermolecular (between molecule) forces of attraction calledHYDROGEN BONDS. As a result of H-bonding, water has an unusuallyhigh melting point and boiling point
compared with similar molecules like H2S. Water has a very highspecific heat, so it heats up and cools down much more slowly
than most materials. This value is found on Table B. Watersolutions that contain ions are capable of conducting electricity.The substances
that form the ions in solution are called electrolytes. 2. Wateris able to make solutions with many substances. Solutions areALWAYS HOMOGENEOUS MIXTURES. Water will dissolve many ioniccompounds. Water will dissolve molecular substances if they arealso polar. This reminds us of the like dissolves like principle.Acids dissolve in water to form H+ ions. (This includes organicacids: R-COOH) Bases dissolve in water to form OH- ions. (This doesNOT include alcohols: R-OH)
3. Ionic compounds may be either soluble or insoluble in water.Use Table F to decide!
4. Solubility describes how much of a particular solute willdissolve in a set amount of water at a certain temperature. UseTable G. The amount of water used is 100 g. Saturated solutionshold all the solute possible at the temperature chosen for thewater. An increase in temperature of the water usually makes itcapable of dissolving more
solute. The opposite is true for gas solutes like O2 gas, or NH3or SO2 or HCl. An increase in pressure over the solution increasesthe solubility of gas solutes. It does
not affect solubility of solutes that are liquids or solids. 5.Solutions have a lower freezing point and a higher boiling pointthat pure water. This effect becomes larger with more concentratedsolutions. 6. Concentration describes how much solute is dissolvedin a certain amount of water. You should know how to calculate:
o Molarity o % mass o Parts per million
Use Table T and plug and chug.
Water & Solutions Cut from Jan 2007 Jan 2008 Exams
1. A 3.0 M HCl(aq) solution contains a total of
(1) 3.0 grams of HCl per liter of water (2) 3.0 grams of HCl permole of solution (3) 3.0 moles of HCl per liter of solution (4) 3.0moles of HCl per mole of water
2. A dilute, aqueous potassium nitrate solution is
best classified as a (1) homogeneous compound (2) homogeneousmixture (3) heterogeneous compound (4) heterogeneous mixture
3. According to one acid-base theory, a water
molecule acts as an acid when the water molecule (1) accepts anH+ (3) donates an H+ (2) accepts an OH (4) donates an OH
4. An Arrhenius base yields which ion as the only negative ionin an aqueous solution?
(1) hydride ion (3) hydronium ion (2) hydrogen ion (4) hydroxideion
5. Which barium salt is insoluble in water?
(1) BaCO3 (3) Ba(ClO4)2 (2) BaCl2 (4) Ba(NO3)2
6. Which unit can be used to express solution concentration?
(1) J/mol (3) mol/L (2) L/mol (4) mol/s
7. Under which conditions of temperature and pressure is a gasmost soluble in water? (1) high temperature and low pressure (2)high temperature and high pressure (3) low temperature and lowpressure (4) low temperature and high pressure 8. Given theequation representing a system at equilibrium:
At which temperature does this equilibrium exist at 101.3kilopascals?
(1) 0 K (3) 32 K (2) 0C (4) 273C
9. As water is added to a 0.10 M NaCl aqueous solution, theconductivity of the resulting solution (1) decreases because theconcentration of ions decreases (2) decreases, but theconcentration of ions remains the same (3) increases because theconcentration of ions decreases (4) increases, but theconcentration of ions remains the same 10. Which substance is anArrhenius acid?
(1) Ba(OH)2 (3) H3PO4 (2) CH3COOCH3 (4) NaCl
11. Which compound releases hydroxide ions in an aqueoussolution?
(1) CH3COOH (3) HCl (2) CH3OH (4) KOH
12. Which liquid has the highest vapor pressure at 75C?
(1) ethanoic acid (3) propanone (2) ethanol (4) water
13. Which sample of matter is a single substance?
(1) air (3) hydrochloric acid (2) ammonia gas (4) salt water
14. At standard pressure, a certain compound has a low boilingpoint and is insoluble in water. At STP, this compound most likelyexists as
(1) ionic crystals (2) metallic crystals (3) nonpolar molecules(4) polar molecules
15. An unsaturated solution is formed when 80. grams of a saltis dissolved in 100. grams of water at 40.C. This salt could be
(1) KCl (3) NaCl (2) KNO3 (4) NaNO3
16. Which substance, when dissolved in water, forms a solutionthat conducts an electric current?
(1) C2H5OH (3) C12H22O11 (2) C6H12O6 (4) CH3COOH
17. Compared to a 2.0 M aqueous solution of NaCl at 1atmosphere, a 3.0 M aqueous solution of NaCl at 1 atmosphere has a(1) lower boiling point and a higher freezing
point (2) lower boiling point and a lower freezing
point (3) higher boiling point and a higher freezing
point (4) higher boiling point and a lower freezing
18. A student prepares four aqueous solutions, each with adifferent solute. The mass of each dissolved solute is shown in thetable below.
Which solution is saturated?
(1) 1 (3) 3 (2) 2 (4) 4
Base your answers to question 19 on the information below.
The equilibrium equation below is related to the manufacture ofa bleaching solution. In this equation, Cl(aq) means that chlorideions are surrounded by water molecules.
19. Draw two water molecules in the box, showing the correctorientation of each water molecule toward the chloride ion. 
Base your answers to questions 20 through 22 on the informationbelow.
Scientists who study aquatic ecosystems are often interested inthe concentration of dissolved oxygen in water. Oxygen, O2, has avery low solubility in water, and therefore its solubility isusually expressed in units of milligrams per 1000. grams of waterat 1.0 atmosphere. The graph below shows a solubility curve ofoxygen in water.
20. A student determines that 8.2 milligrams of oxygen isdissolved in a 1000.-gram sample of water at 15C and 1.0atmosphere. In terms of saturation, what type of solution is thissample?  ___________________ 21. Explain, in terms of molecularpolarity, why oxygen gas has low solubility in water. Your responsemust include both oxygen and water.  22. An aqueous solution has0.0070 gram of oxygen dissolved in 1000. grams of water. Calculatethe dissolved oxygen concentration of this solution in parts permillion. Your response must include both a correct numerical setupand the calculated result. 
Base your answers to questions 23 and 24 on the informationbelow.
A solution is made by completely dissolving 90. grams of KNO3(s)in 100. grams of water in a beaker. The temperature of thissolution is 65C.
23. Describe the effect on the solubility of KNO3(s) in thissolution when the pressure on the solution increases.  24.Determine the total mass of KNO3(s) that settles to the bottom ofthe beaker when the original solution is cooled to 15C.  Baseyour answers to questions 25 through 27 on the informationbelow.
The compound 1,2-ethanediol can be mixed with water. Thismixture is added to automobile radiators as an engine coolant. Thecooling system of a small van contains 6690 grams of1,2-ethanediol. Some properties of water and 1,2-ethanediol aregiven in the table below.
25. Identify the class of organic compounds to which1,2-ethanediol belongs.  _________________________ 26. State, interms of molecular polarity, why 1,2-ethanediol is soluble inwater.  27. Calculate the total number of moles of 1,2-ethanediol in the small vans cooling system. Your response mustinclude both a correct numerical setup and the calculated result.
28. An aqueous solution contains 300. parts per million of KOH.Determine the number of grams of KOH present in 1000. grams of thissolution.  29. A liquid boils when the vapor pressure of theliquid equals the atmospheric pressure on the surface of theliquid. Using Reference Table H, determine the boiling point ofwater when the atmospheric pressure is 90. kPa. 
Topic 8: Acids & Bases Outline
1. An electrolyte is a substance which when dissolved in waterforms a solution capable of conducting an electric current. Theability of a solution to conduct an electric current depends uponthe concentration of ions present. Ionic compounds are conductorsof electricity when melted OR dissolved in water. Under these
circumstances, the charged particles (ions in this case) arefree to move (mobile). There are 3 categories of electrolytes:acids, bases and salts. Arrhenius theory states that an acid is ansubstance that dissolves in water to produce H+
(H3O+) ions (called hydronium ions on Table E). Arrhenius theorystates that a base is an substance that dissolves in water toproduce OH- ions
(called hydroxide ions on Table E). A salt is any ionic compoundproducing a positive ion other than H+ and a negative ion otherthan
OH-. Common acid and base names and formulas are given on TablesK and L. You should be able to sort compounds as acids, bases orsalts, given their chemical formulas
2. Properties of many acids and bases can be explained by theArrhenius theory. Arrhenius acids and bases are electrolytes. Acidproperties include sour taste, less than 7 pH, ability toneutralize bases, and ability to affect
indicator colors. These properties are due to the H+ ion. Baseproperties include bitter taste, greater than 7 pH, ability toneutralize acids, and ability to
affect indicator colors. These properties are due to the OH-ion. When given properties, you should be able to identifysubstances as Arrhenius acids or Arrhenius
3. The acidity or alkalinity of a solution can be measured byits pH value. For every change in pH of one unit, the aciditychanges by a factor of 10. A pH 4 solution is 10
times more acidic that a pH 5 solution. A pH 4 solution is 100times more acidic that a pH 6 solution.
You should be able to identify solutions as acid, base, orneutral based upon the pH. Neutral is a pH of 7.
4. The relative level of acidity or alkalinity of a solution canbe shown by using indicators.
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