Sat chemistry notes

156
CHEMISTRY Dr. D. Bampilis

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SAT Chemistry Notes

Transcript of Sat chemistry notes

  • SAT CHEMISTRY Dr. D. Bampilis 1

    CHEMISTRY

    Dr. D. Bampilis

  • SAT CHEMISTRY Dr. D. Bampilis 2

  • SAT CHEMISTRY Dr. D. Bampilis 3

    SAT Chemistry

    Introduction to Chemistry

    You have to know:

    md

    V

    Matter

    Mixture Substance

    Homogenous Heterogeneous Compound Element

    Chemical - Physical Properties

    E = m.c2

    SI

    Prefixes T/G/M/k d/c/m//n/p

    T = + 273

    = C + 273

    Significant figures

    The result of a calculation cannot be more accurate than the least accurate number in the

    calculation.

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    1. Atomic structure

    To describe the location of electrons, we use quantum numbers. Quantum numbers are basically used to describe certain aspects of the locations of electrons. For example, the quantum numbers n, l, and ml describe the position of the electron with respect to the nucleus, the shape of the orbital, and its special orientation, while the quantum number ms describes the direction of the electrons spin within a given orbital. Below are the four quantum numbers, showing how they are depicted and what aspects of electrons they describe.

    Principal quantum number (n)

    Has positive values of 1, 2, 3, etc. As n increases, the orbital becomes larger-this means that the electron has a higher energy level and is less tightly bound to the nucleus.

    Second quantum number or azimuthal quantum number (l)

    Has values from 0 to n - 1. This defines the shape of the orbital, and the value of l is designated by the letters s, p, d, and f, which correspond to values for l of 0, 1, 2, and 3. In other words, if the value of l is 0, it is expressed as s; if l = 1 = p, l = 2 = d, and l = 3 = f.

    Magnetic quantum number (ml)

    Determines the orientation of the orbital in space relative to the other orbitals in the atom. This quantum number has values from -l through 0 to +l.

    Spin quantum number (ms)

    Specifies the value for the spin and is either +1/2 or -1/2. No more than two electrons can occupy any one orbital. In order for two electrons to occupy the same orbital, they must have opposite spins.

    Orbitals that have the same principal quantum number, n, are part of the same electron shell. For example, orbitals that have n = 2 are said to be in the second shell. When orbitals have the same n and l, they are in the same subshell; so orbitals that have n = 2 and l = 3 are said to be 2f orbitals, in the 2f subshell. Finally, you should keep in mind that according to the Pauli exclusion principle, no two electrons in an atom can have the same set of four quantum numbers. This means no atomic orbital can contain more than two electrons, and if the orbital does contain two electrons, they must be of opposite spin.

    You have to know:

    Proton - Neutron - Electron

    Atomic number (Z) - Isotopes

    Mass number (A)

    Bohr model - atomic spectra

    Lymar - Balmer - Paschen

    Mass spectroscopy

    (vaporizer - ionizer - accelerator - deflector - detector - recorder)

    The wave - mechanical model

    Max Planck

    Louis de Broglie

    Heisenberg

    Quantum numbers

    Principal quantum number (n): K:1, L:2, M:3, N:4

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    Angular momentum (l) quantum number l = 0n-1

    l: 0 (s),1(p),2(d),3(f)

    Magnetic quantum number (ml) ml = -l, .,0, . +l

    Spin quantum number (ms)

    Orbitals

    Pauli Exclusion Principle

    Hunds rule

    Aufbau Principle

    Electron configuration

    Exercises

    1. An element consists of three isotopes in the relative abundance given below.

    What is the atomic mass of this element?

    30.00% = 40.00 amu

    50.00% = 41.00 amu

    20.00% = 42.00 amu

    () 40.90 () 41.00 (C) 41.9 (D) 42.20 () 42.90

    2. The total number of electrons that can be accommodated in the fourth

    principal energy level is:

    () 2 () 8 (C) 18 (D) 32 () 50

    3. If the set of quantum numbers n = 3, l = 1, ml = 0, ms = +1/2 represents the last

    electron to be added to complete the ground state electron configuration of an

    element, which one of the following could be the symbol for the element?

    (A) Na (B) Si (C) Th (D) V (E) Zn

    4. Which element has the electron configuration:1s22s22p63s23p64s23d4?

    () Cr () n (C) o (D) S () Se

    5. Which of the following elements has electrons in f orbitals?

    () Ar () O (C) S (D) Ti () U

    6. Which of the following elements has the electron configuration: 1s22s22p63s23p4?

    () Ar () O (C) S (D) Ti () U

    7. Which of the following elements has the same number of electrons as Ca2+?

    () Ar () O (C) S (D) Ti () U

    8. Which gas-phase atom in its ground state could have an electron with the

    quantum numbers: n = 3, l = 2, ml = 0, ms = -1/2?

    (A) Na (B) Mg (C) P (D) Ti

    9. Which set of quantum numbers could represent an electron in a 5f orbital?

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    (A) l = 4, ml = 2 (B) l = 2, ml = -3 (C) l = 3, ml = 4 (D) l = 3, ml = 0

    10. Which set of quantum numbers (n, l, ml, ms) is possible for the outermost electron in a

    strontium atom in its ground state?

    (A) 5, 0, 0, -1/2 (B) 5, 0, 1, 1/2 (C) 5, 1, 0, 1/2 (D) 5, 1, 1, -1/2

    11. Which quantum numbers represent the orbitals being filled in the ground state for the

    elements Sc (21) to Zn (30)?

    (A) n = 3, l = 1 (B) n = 3, l = 2 (C) n = 4, l = 1 (D) n = 4, l = 2

    12. Which set of quantum numbers (n, l, ml) is forbidden?

    (A) 3, 2, 0 (B) 3, 1, 1 (C) 2, 0, 0 (D) 1, 1, 0

    13. Which characteristic of an atomic orbital is most closely associated with the magnetic

    quantum number, ml?

    (A) size (B) shape (C) occupancy (D) orientation

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    2. Periodicity

    You have to know:

    Electron configuration and Periodic Table

    s, p, d, f block

    Transition elements - d block

    d: half - filled or full - filled

    Periodic law

    Alkali metals (1) - Alkaline Earth Metals (2) - Halogens (17) - Noble gases (18) - Metals - Nom

    metals - Metalloids (B, Si, Ge, As, Sb, Te)

    Lanthanides - Actinides: f block

    Atomic Radii

    Ionic radius

    Electronegativity

    Electron affinity

    Ionization Energy

    Why is fluorine out of line?

    The incoming electron is going to be closer to the nucleus in fluorine than in any other of these elements, so you would expect a high value of electron affinity. However, because fluorine is such a small atom, you are putting the new electron into a region of space already crowded with electrons and there is a significant amount of repulsion. This repulsion lessens the attraction the incoming electron feels and so lessens the electron affinity.

    A similar reversal of the expected trend happens between oxygen and sulfur in Group 6. The first electron affinity of oxygen (-142 kJ.mol-1) is smaller than that of sulfur (-200 kJ.mol-1) for exactly the same reason that fluorine's is smaller than chlorine's.

    Second electron affinity

    You are only ever likely to meet this with respect to the group 6 elements oxygen and sulfur which both form -2 ions. The second electron affinity is the energy required to add an electron to each ion in 1 mole of gaseous 1- ions to produce 1 mole of gaseous 2- ions. This is more easily seen in symbol terms.

    X(g) + e X2(g)

    It is the energy needed to carry out this change per mole of X.

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    Why is energy needed to do this?

    You are forcing an electron into an already negative ion. It's not going to go in willingly!

    O(g) + e- O-(g) 1st EA = -142 kJ.mol

    -1

    O-(g) + e- O2-(g) 2nd EA = +844 kJ.mol

    -1

    The positive sign shows that you have to put in energy to perform this change. The second electron affinity of oxygen is particularly high because the electron is being forced into a small, very electron-dense space.

    Exercises

    1. Which of the following atoms would have the largest second ionization energy?

    () Mg () Cl (C) S (D) Ca () a

    2. Order the elements S, CI, and F in terms of increasing atomic radii.

    () S, CI, F () CI, F, S (C) F, S, CI (D) F, CI, S () S, F, CI

    3. Which of the following elements is the LEAST chemically reactive?

    () Ar () O (C) S (D) Ti () U

    4. For elements in the left-most column of the periodic table, properties that have increasing

    values as the atomic number increases include which of the following?

    I. Ionization energy (potential) II. Atomic radius III. Atomic mass

    (A) I only (B) III only (C) I and II only (D) II and III only (E) I, II, and III

    5. Removing an electron from sodium is an ..... process and removing an electron

    from fluorine is an ..... process.

    () Endothermic, exothermic () Exothermic, endothermic

    (C) Endothermic, endothermic (D) Exothermic, exothermic

    () More information is needed

    For questions 6 - 10: Statement I BECAUSE Statement II

    6. The second ionization energy of is higher than that of Be.

    BECAUSE

    The second electron to be removed from and Be comes from the same principal

    energy level.

    7. Oxygen has a smaller first ionization energy than fluorine.

    BECAUSE

    Oxygen has a higher eff value than does fluorine.

    8. Hydrogen has a lower ionization energy than does helium.

    BECAUSE

    Hydrogen bonds with halogens to form polar covalent bonds.

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    9. Potassium has a lower first ionization energy than lithium has.

    BECAUSE

    Potassium has more protons in its nucleus than lithium has.

    10. An element that has the electron configuration 1s22s22p63s23p63d34s2 is a transition

    element.

    BECAUSE

    In atoms of transition elements, the 1s, 2s, 2p, 3s and 3p orbitals are completely filled in the

    ground state.

    Atomic structure & periodicity For questions 1 - 5:

    () Alkali metals () Alkaline Earth metals (C) Metalloids

    (D) Halogens (E) Rare earth metals

    1. Which of the following is used primarily in semiconductors?

    2. Which occur as diatomics?

    3. Which make oxides with the formula X2O?

    4. Which have large electronegativity values?

    5. Which have small ionization energies?

    For questions 6 - 9:

    () Na+ (B) Al (C) F (D) Ti (E) B

    6. Which has seven valence electrons?

    7. Which has an electron configuration 1s22s22p63s23p1?

    8. Which has the same electron configuration as the neon atom?

    9. Which has valence electrons in the d orbitals?

    For questions 10 - 14:

    () Alkali metals (B) Alkaline Earth metals (C) Noble gases

    (D) Halogens (E) Transition metals

    10. Which is the most unreactive family of elements?

    11. Which form negative ions in an ionic bond?

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    12. Which consist of atoms that have valence electrons in a d subshell?

    13. Which exist as diatomic molecules at room temperature?

    14. Which group possesses the lowest first ionization energy in their respective period?

    For questions 15 - 17:

    () Bohr model (B) De Broglies wave hypothesis

    (C) Heisenbergs uncertainty principle (D) Quantum theory (E) Atomic theory

    15. Which principle provides that all matter may be considered a wave?

    16. What views electrons in true orbits around the nucleus?

    17. What considers that one cannot know position and velocity of an electron at the same

    moment?

    For questions 18 - 32: Statement I BECAUSE Statement II

    18. The metalloids have similar characteristics

    BECAUSE

    Their valence shells have the same configuration

    19. Metals are good conductors of heat and electricity

    BECAUSE

    The positive nuclei are surrounded by a sea of mobile electrons

    20. Elements in a group have similar properties

    BECAUSE

    Their valence shells have the same energy

    21. The first ionization energy for an atom is greater than the second ionization energy

    BECAUSE

    The closer an electron is to the nucleus, the more difficult it is to remove

    22. Sodium has a smaller atomic radius than chlorine

    BECAUSE

    A sodium atom does not have as many valence electrons as a chlorine atom does.

    23. Carbons electric configuration is 1s22s22p2 rather than 1s22s23s2

    BECAUSE

    3s electrons are lower in energy than 2p electrons

    24. The properties of phosphorus should be closer to those of sulfur than to those of

    nitrogen

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    BECAUSE

    Phosphorus and nitrogen are in the same row of the periodic table

    25. The halogens, in group VIIA, all form stable diatomic molecules

    BECAUSE

    They each need one electron to fill their outer shells

    26. Metals are good conductors of electricity

    BECAUSE

    They are held together by ionic bonds

    27. Two electrons in the 2s subshell must have opposite spins

    BECAUSE

    The Pauli exclusion principle states that no two electrons in the same atom can have

    identical quantum numbers

    28. 40Ca is a neutral atom

    BECAUSE

    It has the same number of protons and neutrons

    29. The most important factor in determining the chemical properties of an element is the

    number of electrons in the outermost shell

    BECAUSE

    The number of electrons in the outermost shell determines the bonding characteristics of

    the element

    30. Iron is an element

    BECAUSE

    It cannot be broken into smaller units and retains its physical and chemical properties

    31. An element (X) with an atomic number of 16 has 14 electrons in X2+

    BECAUSE

    It has gained two electrons

    32. Atomic radii increase down a group

    BECAUSE

    The higher the atomic number within a group, the smaller the atom

    33. The element with atomic number 32 describes:

    () A metal (B) A non-metal (C) A metalloid (D) A halogen (E) A noble gas

    34. How many neutrons are probably in the nucleus of an element of atomic weight 197?

    () 43 (B) 79 (C) 83 (D) 100 (E) 118

    35. The transition metals are characterized by:

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    (A) completely filled d subshells (B) completely filled f subshells

    (C) partially filled d subshells (D) partially filled f subshells

    (E) both (a) and (c) are correct

    36. Neutral atoms of F (fluorine) have the same number of electrons as:

    (A) B3- (B) N+ (C) Ne- (D) Na- (E) Mg3+

    37. Which of the elements in Group 1A of the periodic table has the greatest metallic

    character?

    (A) H (B) Li (C) Na (D) K (E) Rb

    38. The ionization energy of an element is:

    (A) a measure of its mass

    (B) the energy required to remove an electron from the element in its gaseous state

    (C) the energy released by the element in forming an ionic bond

    (D) the energy released by the element upon receiving an additional electron

    (E) none of the above

    39. Elements in a row have the same:

    (A) Atomic weight (B) Maximum azimuthal quantum number(l)

    (C) Maximum principal quantum number (n) (D) Valence electron structure

    (E) Atomic number

    40. Which of the following has the largest radius?

    (A) Sr (B) P (C) Mg (D) Al3+ (E) Mg2+

    41. Which of the following elements has the lowest electronegativity?

    (A) Cesium (B) Strontium (C) Calcium (D) Barium (E) Potassium

    42. Which of the following is biggest in size?

    (A) Ca (B) Ca+ (C) Ca2+ (D) Ca- (E) Ca2-

    43. The order of the elements in the periodic table is based on:

    (A) the number of neutrons (B) the radius of the atom (C) the atomic number

    (D) the atomic weight (E) the number of oxidation states

    44. The elements within each column of the Periodic Table:

    (A) have similar valence electron configurations (B) have similar atomic radii

    (C) have the same principal quantum number (D) will react to form stable elements

    (E) have no similar chemical properties

    45. Which of the following has the highest 1st ionization energy?

    (A) Ga (B) Ba (C) Ru (D) F (E) N

    46. Which of the following has the lowest electronegativity?

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    (A) Ca (B) Cl (C) Cs (D) P (E) Zn

    47. Which element has the greatest electronegativity?

    (A) Chlorine (B) Oxygen (C) Sulfur (D) Phosphorus (E) Fluorine

    48. Transition metal compounds generally exhibit bright colors because:

    (A) The electrons in the partially filled d orbitals are easily promoted to excited states

    (B) The metals become complexed with water

    (C) The metals conduct electricity, producing colored light

    (D) The electrons in the d orbitals emit energy as they relax

    (E) Their valence electrons cause them to bind to other metals

    49. Which of the following is a non metal?

    (A) Fr (B) Pd (C) I (D) Sc (E) Sr

    50. Which of the following has the greatest affinity for electrons?

    (A) F (B) Cl (C) Br (D) K (E) C

    51. Which of the following is the most electronegative element?

    (A) He (B) I (C) N (D) O (E) C

    52. Which of the following is not a property of Group IA elements?

    (A) Low ionization energies (B) Low electronegativities (C) High melting points

    (D) Metallic bonding (E) Electrical conductivity

    53. Arrange the following elements in order of decreasing nonmetallic character: Ge, Sn, Pb,

    Si.

    (A) Pb, Sn, Ge, Si (B) Ge, Sn, Pb, Si (C) Si, Ge, Sn, Pb

    (D) They all have equal nonmetallic character since they are all in the same column of the

    Periodic Table

    (E) None of the above

    54. Electron affinity is defined as:

    (A) the change in energy when a gaseous atom in its ground state gains an electron

    (B) the pull an atom has on the electrons in a chemical bond

    (C) the energy required to remove a valence electron from a neutral gaseous atom in its

    ground state

    (D) the energy difference between an electron in its ground state and its excited state

    (E) none of the above

    55. Which of the following is an incorrect association?

    (A) Mendeleev-periodic table (B) Faraday-electrolytic cells

    (C) Millikan-charge of electrons (D) Rutherford-photoelectric effect

    (E) They are all correct

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    56. Members of group 1 have similar reactivity because they have:

    (A) the same number of protons (B) the same number of electrons

    (C) similar outer shell configurations

    (D) valence electrons with the same quantum numbers

    (E) the same number of neutron

    57. Boron found in nature has an atomic weight of 10.811 and is made up of the isotopes 10B

    (mass 10.013 amu) and 11B (mass 11.0093 amu). What percentage of naturally occurring

    boron is made up of 10B and 11B respectively?

    (A) 30:70 (B) 25:75 (C) 20:80 (D) 15:85

    58. The modern periodic table is ordered on the basis of:

    (A) atomic mass (B) atomic radius (C) atomic charge

    (D) atomic number (E) number of neutron

    59. The electron configuration 1s22s22p63s23p64s23d7 represents an atom of the element:

    (A) Br (B) Co (C) Cd (D) Ga (E) Mg

    60. A neutral atom whose electron configuration is 1s22s22p63s23p64s23d104p65s24d105p6 is:

    (A) Highly reactive (B) A noble gas (C) A positively charged ion

    (D) A transition metal (E) A lanthanide element

  • SAT CHEMISTRY Dr. D. Bampilis 15

    3. Bonding

    You have to know:

    Ionic 1.7

    Covalent (polar - non polar) - Lewis structure

    Exceptions to the octet rules: BeH2, BF3, BH3, PCl5, SCl6

    Coordinate Covalent bonds

    Metallic

    Intermolecular forces (Van der Walls - SAT)

    Dipole - Dipole attraction

    London Dispersion Forces (Van der Walls - IB)

    Hydrogen Bonds

    Molecule - ion attraction

    Resonance structure

    VSEPR polar - non polar molecules

    Hybridization

    Sigma - pi bonds

    Properties

    Ionic substances Molecular crystals and liquids

    Solid: dont conduct electric current

    liquid: conduct electric current

    Solid or liquid: dont conduct electric current

    Melting Point Boiling Point: high

    Low volatilities and vapor pressures

    Melting Point Boiling Point: Low

    Relative volatile

    Brittle solids Soft - waxy (solids)

    Soluble in water (with exceptions) Large amount of energy to decompose the

    substance

    Exercises

    For questions 1 - 4:

    (A) an ionic substance (B) a polar covalent substance (C) a nonpolar covalent substance

    (D) an amorphous substance (E) a metallic network

    1. KCl(s) is:

    2. HCl(g) is:

    3. CH4(g) is:

    4. Li(s) is:

    For questions 5 - 8:

    (A) hydrogen bond (B) ionic bond (C) polar covalent bond (D) pure covalent bond

    (E) metallic bond

  • SAT CHEMISTRY Dr. D. Bampilis 16

    5. The type of bond between atoms of potassium and chloride in a crystal of potassium

    chloride is:

    6. The type of bond between the atoms in a nitrogen molecule is:

    7. The type of bond between atoms in a molecule of CO2 (electronegativity difference = ~1)

    is:

    8. The type of bond between atoms of calcium in a crystal of calcium is:

    For questions 9 - 11:

    (A) zero (B) one (C) two (D) three (E) four

    9. The number of bonds predicted for O2.

    10. The number of bonds predicted for N2.

    11. The number of bonds predicted for H2.

    For questions 12 - 15:

    (A) Linear geometry (B) Bent geometry (C) Tetrahedral geometry

    (D) Pyramidal geometry (E) Equilateral triangle geometry

    12. NH3 has a:

    13. H2O has a:

    14. BeF2 has a:

    15. CH4 has a:

    For questions 16 - 18:

    (A) BeF2 (B) NH3 (C) CH4 (D) CH2CH2 (E) CCl4

    16. This species has sp2 hybrid orbitals.

    17. This species has sp hybrid orbitals.

    18. This species contains a pi bond.

    For questions 19 - 22:

    (A) hydrogen bonding (B) ionic bonding (C) metallic bonding

    (D) nonpolar covalent bonding (E) polar covalent bonding

    19. This holds a sample of barium iodide, BaI2, together.

  • SAT CHEMISTRY Dr. D. Bampilis 17

    20. This allows many solids to conduct electricity.

    21. This attracts atoms of hydrogen to each other in a H2 molecule.

    22. This is responsible for the relatively high boiling point of water.

    For questions 23 - 32: Statement I BECAUSE Statement II

    23. Nonmetallic atoms of the same element combine covalently

    BECAUSE

    The two elements have the same electronegativities

    24. A nonpolar molecule can have polar bonds

    BECAUSE

    Polar bonds can be symmetrically arranged in a molecule so that there are no net poles

    25. The bond in an O2 molecule is considered to be nonpolar

    BECAUSE

    The oxygen atoms in an O2 molecule share the bonding electrons equally

    26. An ionic solid is a good conductor of electricity

    BECAUSE

    An ionic solid is composed of positive and negative ions joined together by electrostatic

    forces

    27. The hybrid orbitals of carbon in acetylene are believed to be the sp form

    BECAUSE

    Acetylene is a linear compound with a triple bond between the carbons

    28. Atom A with 7 valence electrons forms AB2 with atom B with two valence electrons

    BECAUSE

    B donates its electrons to fill the outer shell of A

    29. Water is a polar substance

    BECAUSE

    The bonding electrons in water are shared equally

    30. He2 is not known to commonly form

    BECAUSE

    He is lighter than air

    31. CCl4 is a nonpolar molecule

    BECAUSE

    The dipole moments in CCl4 cancel each other out

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    32. One of the most important factors in determining the chemical properties of an element

    is the number of electrons in its outermost shell

    BECAUSE

    The number of electrons in the outer shell determines the bonding characteristics of an

    element

    33. An sp2 configuration is represented by which orientation?

    (A) Tetrahedral (B) Planar (C) Linear (D) Trigonal planar (E) Square

    34. When the electrons are shared unequally by two atoms, the bond is said to be:

    (A) covalent (B) polar covalent (C) coordinate covalent (D) ionic (E) metallic

    35. Which of the following contains a coordinate covalent bond?

    (A) HCl () H2O (C) H2 (D) H3O+ (E) NaCl

    36. Which of the following elements can form bonds with sp3 hybridization?

    (A) Sodium (B) Nitrogen (C) Carbon (D) Oxygen (E) Fluorine

    37. A triple bond may be best described as:

    (A) two sigma bonds and one pi bond (B) one sigma bond and two pi bonds

    (C) two sigma bonds and two pi bonds (D) three sigma bonds

    (E) three pi bonds

    38. Molecules of sodium chloride:

    (A) display ionic bonding (B) display polar covalent bonding

    (C) are polar (D) dissociate in water solution (E) do not exist

    39. Which of the following molecules is polar?

    (A) BH3 (B) NF3 (C) C2H6 (D) SF6 (E) CCl4

    40. Which of the following molecules has a trigonal pyramidal geometry?

    (A) BH3 (B) H2O (C) CH4 (D) NH3 (E) AlCl3

    41. The shape of a PCl3 molecule is described as:

    (A) bent (B) trigonal pyramidal (C) linear (D) trigonal planar (E) tetrahedral

    42. The structure of BeCl2 can best be described as:

    (A) linear (B) bent (C) trigonal (D) tetrahedral (E) square

    43. All of the following have covalent bonds EXCEPT:

    (A) HCl (B) CCl4 (C) H2O (D) CsF (E) CO2

    44. The complete loss of an electron of one atom to another atom with the consequent

    formation of electrostatic charges is said to be:

    (A) A covalent bond (B) A polar covalent bond (C) An ionic bond

  • SAT CHEMISTRY Dr. D. Bampilis 19

    (D) A coordinate covalent bond (E) A pi bond between p orbitals

    45. Which molecule is incorrectly matched with the molecular geometry?

    Molecule Molecular geometry

    (A) SF6 Octahedral

    (B) CH4 Tetrahedral

    (C) SO3 trigonal planar

    (D) SeCl4 tetrahedral

    (E) PH3 trigonal pyramidal

    46. From their electron configurations, one can predict that the geometric configuration for

    which of the following molecules is NOT correct?

    (A) PF3 trigonal planar (B) CF4 tetrahedral (C) CHCl3 irregular tetrahedron

    (D) OF2 bent (v-shaped) (E) HF linear

    47. Which numbered response lists all the molecules below that exhibit

    resonance and none that do not?

    . AsF5 . O3 III. SO2

    (A) I only (B) II only (C) and III (D) III and IV () , , and III

    48. Which one of the following molecules is octahedral?

    () BeCl2 () SeF6 (C) BF3 (D) PF5 () CF4

    49. The sulfur hexafluoride molecule is nonpolar and contains no lone

    (unshared) electron pairs on the sulfur atom. Which answer choice lists all of the

    bond angles contained in sulfur hexafluoride?

    () 120o (B) 180o (C) 90o and 180o (D)90o, 120o, and 180o ()109.5o

    For questions 50 - 54: Statement I BECAUSE Statement II

    50. Molecules that contain a polar bond are not necessarily polar compounds.

    BECAUSE

    If polar bonds in a molecule are symmetrically arrange d, then their polarities will

    cancel and they will be nonpolar.

    51. The 3 molecule is more polar than the NF3 molecule.

    BECAUSE

    Fluorine atoms are larger than hydrogen atoms.

    52. Ice, unlike most substances, is denser than water in the liquid phase.

    BECAUSE

    In water, hydrogen bonds can form between the positively charged atom on one

    water molecule and the slightly negatively charged atom on a nearby water

    molecule.

  • SAT CHEMISTRY Dr. D. Bampilis 20

    53. Diamond has a high melting point.

    BECAUSE

    In a diamond crystal, the carbon atoms are held in place by ionic bonds.

    54. A molecule of silicon tetrachloride, SiCl4 , is nonpolar.

    BECAUSE

    The four bonds in SiCl4 are identical and the molecule has a tetrahedral structure.

  • SAT CHEMISTRY Dr. D. Bampilis 21

    4. Gases and Gas Laws

    Grahams Law of Diffusion and Effusion Grahams law states that the rates of effusion of two gases are inversely proportional to the square roots of their molar masses at the same temperature and pressure:

    Effusion is the term used to describe the passage of a gas through a tiny orifice into an evacuated chamber, as shown in the figure below.

    The rate of effusion measures the speed at which the gas travels through the tiny hole into a vacuum. Another term to remember for the test is diffusion. Diffusion is the term used to describe the spread of a gas throughout a space or throughout a second substance.

    You have to know:

    Kinetic Molecular Theory

    Maxwell - Boltzmann distribution

    Grahams law of diffusion (effusion)

    B

    A

    MRate A

    Rate B M

    Charless law V

    kT

    Boyles law P V = k

    Gay - Lussacs law P

    kT

    Combined gas law 1 1 2 2

    1 2

    PV P V

    T T

    Daltons law of Partical Pressures

    Ptot = P1 + P2 + P3

    11 tot

    tot

    nP P

    n

    Ideal gas law: P V = nRT

    (real gases at low pressure and high temperature)

    Two devices to measure pressure:

    Mercury Barometer (Eudiometer)

    Manometer

  • SAT CHEMISTRY Dr. D. Bampilis 22

    Exercises

    1. Which of the following gases would be the densest at standard temperature and

    pressure?

    () Helium () Argon (C) Carbon dioxide (D) Xenon () Nitrogen

    2. A gas diffuses one-third as fast as O2 at 100oC. This gas could be:

    (A) He (M = 4) (B) C2H5F (M = 48) (C) C7H12 (M = 96) (D) C5F12 (M = 288)

    3. Ar and He are both gases at room temperature. How do the average molecular velocities

    (V) of their atoms compare at this temperature?

    (A) VHe = 10VAr (B) VAr = 10VHe (C) VHe = 3VAr (D) VAr = 3VHe

    4. Moist air is less dense than dry air at the same temperature and barometric pressure.

    Which is the best explanation for this observation?

    (A) H2O is a polar molecule but N2 and O2 are not.

    (B) H2O has a higher boiling point than N2 or O2.

    (C) H2O has a lower molar mass than N2 or O2.

    (D) H2O has a higher heat capacity than N2 or O2.

    5. Two samples of gas, one of argon and one of helium, have the same pressure,

    temperature and volume. Which statement is true assuming both gases behave ideally?

    (A) The helium sample contains more atoms than the argon sample and the helium atoms

    have a higher average speed.

    (B) The two samples have the same number of atoms but the helium atoms have a higher

    average speed.

    (C) The two samples have the same number of atoms and both types of atoms have the

    same average speed.

    (D) The two samples have the same number of atoms but the argon atoms have a higher

    average speed.

    6. Which of the following best illustrates a graph of pressure versus volume for a

    gas at constant temperature?

  • SAT CHEMISTRY Dr. D. Bampilis 23

    7. The bulb of the open-end manometer shown below contains a gas. True statements about

    this system include which of the following?

    I. Only atmospheric pressure is exerted on the exposed mercury surface in the right side of

    the tube.

    II. The gas pressure is greater than atmospheric pressure.

    III. The difference in the height, h, of mercury levels is equal to the pressure of the gas.

    (A) II only (B) III only (C) I and II only (D) I and III only (E) I, II, III

    For question 8: Statement I BECAUSE Statement II

    8. Statement I: At the same temperature and pressure, 1 L of hydrogen gas and 1 L of neon

    gas have the same mass.

    BECAUSE

    Statement II: Equal volumes of ideal gases at the same temperature and pressure contain the

    same number of moles.

  • SAT CHEMISTRY Dr. D. Bampilis 24

    5. Stoichiometry

    You have to know:

    Mole - Avogadros number

    Avogadros law n

    kV

    Molar mass

    r m A

    m V V Nn n (STP) n

    M V 22,4 N

    m M Md (STP)

    V Vm 22,4

    mP V RT P M d RT

    M

    Balancing chemical equations

    Mole ratio

    Limiting and excess reagents

    Percent yield of a product

    Exercises

    For questions 1 - 3:

    (A) N2O5 (B) N2O3 (C) NO2 (D) NO (E) N2O

    1. What is the empirical formula for a compound containing 63.8% N and 36.2% O?

    2. What is the empirical formula for a compound containing 36.7% N and 63.3% O?

    3. What is the empirical formula for a compound containing 25.9% N and 74.1% O?

    For questions 4 - 6:

    (A) 2.294 (B) 36.51 (C) 1.409 (D) 25.3 (E) 2.513

    4. For 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g), if you begin with 16.00 g ammonia and excess

    oxygen, how many grams of water will be obtained?

    5. For 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g), if you begin with 66.00 g ammonia and 54.00 g

    oxygen, how many grams of water will be obtained?

    6. For 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g), how many moles of NH3 are needed to produce

    2.513 moles of NO?

    For questions 7 - 10:

    (A) 1.807 x 10-24 (B) 3.476 x 10-2 (C) 1.171 x 10-2 (D) 1.204 x 1024 (E) 2.414 x 10-1

  • SAT CHEMISTRY Dr. D. Bampilis 25

    7. How many phosphine molecules are in two moles of phosphine?

    8. How many moles of CO2 are in 1.53 g CO2?

    9. How many atoms are in one mole of water?

    10. How many moles are in 4.35 grams of water?

    11. When the following is balanced, C4H10 + O2 CO2 + H2O, what is the coefficient of CO2?

    (A) 2 (B) 4 (C) 8 (D) 10 (E) 13

    12. What is the approximate percentage composition by mass of the element oxygen in the

    compound HClO4?

    (A) 16% (B) 35% (C) 50% (D) 64% (E) 75%

    13. When the following equation is balanced, how many moles of NF3 would be required to

    react completely with 6 moles of H2O?

    ___NF3(g) + ___H2O(g) ___HF(g) + ___NO(g) + ___NO2(g)

    (A) 0.5 mole (B) 1 mole (C) 2 moles (D) 3 moles (E) 4 moles

    14. For the following equation, Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g), when 3.0 mol Fe2O3 is

    allowed to completely react with 56 g CO, approximately how many moles of iron, Fe, are

    produced?

    (A) 0.7 (B) 1.3 (C) 2.0 (D) 2.7 (E) 6.0

    15. What is the percent by mass of silicon in a sample of SiO2?

    (A) 21% (B) 33% (C) 47% (D) 54% (E) 78%

    16. When the following equation is balanced, ___PH3 + ___O2 ___P2O5 + ___H2O, what is

    the coefficient of H2O?

    (A) 1 (B) 2 (C) 3 (D) 4 (E) 5

    17. What are the products of the following reaction: H2SO4(aq) + Ba(OH)2(aq) ?

    (A) O2 (B) BaSO4 (C) O2 and BaSO4 (D) O2 and BaSO4 (E) H2O and BaSO4

    18. For the equation, 2Mg(s) + O2(g) 2MgO(s), if 48.6 g Mg is placed in a container with

    64.0 g O2 and the reaction is allowed to go to completion, what is the mass of MgO(s)

    produced?

    (A) 15.4 g (B) 32.0 g (C) 80.6 g (D) 96.3 g (E) 112 g

    19. For the equation, 2NO(g) + 2H2(g) N2(g) + 2H2O(g), which of the following is true?

    (A) If 1 mole of H2 is consumed, 0.5 moles of N2 is produced

    (B) If 1 mole of H2 is consumed, 0.5 moles of H2O is produced

    (C) If 0.5 moles of H2 is consumed, 1 mole of N2 is produced

    (D) If 0.5 moles of H2 is consumed, 1 mole of NO is produced

  • SAT CHEMISTRY Dr. D. Bampilis 26

    (E) If 0.5 moles of H2 is consumed, 1 mole of H2O is produced

    20. Which of the following expressions is equal to the number of iron (Fe) atoms present in

    10.0 g Fe? (atomic mass of Fe = 55.9)

    (A) 10 x 55.9 x (6.022 x 1023) atoms (B) (6.022 x 1023) / 10 x 55.9 atoms

    (C) 10 x (6.022 x 1023) / 55.9 atoms (D) 55.9 / 10 x (6.022 x 1023) atoms

    (E) 10 / (55.9 x 6.022 x 1023) atoms

    21. The formula Cr(NH3)5SO4Br represents:

    (A) 4 atoms (B) 8 atoms (C) 12 atoms (D) 23 atoms (E) 27 atoms

    22. What is the molecular formula of a compound made of 25.9% N and 74.1% O?

    (A) NO (B) NO2 (C) N2O (D) N2O5 (E) N2O4

    23. The balanced molar relationship from the reaction H2O2 H2O + O2 is

    (A) 1:1:1 (B) 2:1:1 (C) 1:2:1 (D) 2:2:1 (E) 2:1:2

    24. What volume of H2O is required to produce 5 L O2 by the following equation:

    H2O(g) H2(g) + O2(g)?

    (A) 3 L (B) 5 L (C) 10 L (D) 16 L (E) 14 L

    25. What is the molecular weight of HClO4?

    (A) 52.5 (B) 73.5 (C) 96.5 (D) 100.5 (E) 116.5

    26. Which of the following molecules contains 17 atoms?

    (A) Al2(SO4)3 (B) Al(NO3)3 (C) Ca(HCO2)2 (D) Mg(IO3)2 (E) Two of the above

    27. Twenty liters of NO gas react with excess oxygen. How many liters of NO2 gas are

    produced if the NO gas reacts completely? (2NO + O2 2NO2)

    (A) 5 L (B) 10 L (C) 20 L (D) 40 L (E) 50 L

    28. How much reactant remains if 92 g HNO3 reacts with 24 g LiOH assuming a complete

    reaction?

    (A) 46 g HNO3 (B) 29 g HNO3 (C) 12 g HNO3 (D) 2 g LiOH (E) 12 g LiOH

    29. What is the density, at STP, of a diatomic gas whose gram-formula mass is 80 g/mol?

    (A) 1.9 g/L (B) 2.8 g/L (C) 3.6 g/L (D) 4.3 g/L (E) 5.0 g/L

    30. How many liters of H2 can be produced at STP by the decomposition of 3 mol NH3?

    (A) 4.5 L (B) 27 L (C) 67.2 L (D) 96 L (E) 101 L

    31. How many mol CO2 molecules are represented by 1.8 x 1024 atoms?

    (A) 1 (B) 2 (C) 3 (D) 4 (E) 5

    32. How many grams of Na2SO4 can be produced by reacting 98 g H2SO4 with 40 g NaOH?

  • SAT CHEMISTRY Dr. D. Bampilis 27

    (A) 18 g (B) 36 g (C) 71 g (D)142 g (E) 150 g

    33. What are the missing products of the following reaction?

    NH4Cl + Ca(OH)2 _____ + CaCl2

    (A) N2 (B) NH3 (C) H2O (D) NH3 + N2 (E) NH3 + H2O

    34. How many grams of water can be produced when 8 g of hydrogen react with 8 g oxygen?

    (A) 8 g (B) 9 g (C) 18 g (D) 27 g (E) 30 g

    35. How many atoms are represented in Na2CO310H2O

    (A) 4 (B) 16 (C) 36 (D) 60 (E) 96

    36. What is the density of bromine vapor at STP?

    (A) 2.5 g/L (B) 2.9 g/L (C) 3.6 g/L (D) 4.9 g/L (E) 7.1 g/L

    37. Fill in the missing reactant: NaOH + _____ NaClO2 + H2O

    (A) Cl2 (B) HCl (C) HClO (D) HClO2 (E) HClO3

    38. How many grams of Na are present in 30 g NaOH?

    (A) 10 g (B) 15 g (C) 17 g (D) 20 g (E) 22 g

    39. What is the sum of the coefficients when the following reaction is balanced?

    ___C6H6 + ___O2 ___CO2 + ___H2O?

    (A) 7 (B) 14 (C) 28 (D) 35 (E) 42

    40. How many atoms are represented by the following formula? K3Fe(CN)6

    (A) 6 () 10 (C) 16 (D) 20 (E) 18

    41. Twenty-two grams of CO2 at STP is identical to

    (A) 1 mole of CO2 (B) 6.022 x 1023 atoms (C) 6.022 x 1023 molecules

    (D) 11.2 liters (E) 22.4 liters

    42. What volume does 8.5 g NH3 occupy at STP?

    (A) 2.81 L (B) 5.61 L (C) 11.21 L (D) 22.41 L (E) 44.81 L

    43. What is the formula of a hydrocarbon composed of 86% carbon and 14% hydrogen by

    weight?

    (A) CH4 (B) C2H4 (C) C2H6 (D) C3H8 (E) C4H6

    44. How many grams of CO2 are produced by the complete reaction of 100 g CaCO3 with

    excess HCl?

    (A) 22 g (B) 44 g (C) 79 g (D) 110 g (E) 132 g

    45. 28 mL of nitrogen are reacted with 15 mL of hydrogen. How many milliliters of which gas

    are left unreacted?

  • SAT CHEMISTRY Dr. D. Bampilis 28

    (A) 5 mL H2 (B) 5 mL N2 (C) 7 mL H2 (D) 11 mL N2 (E) 23 mL N2

    46. If 28 mL of nitrogen are reacted with 15 mL of hydrogen, what is the total volume of gas

    present after the reaction has occurred, assuming volumes are additive?

    (A) 11 mL (B) 17 mL (C) 27 mL (D) 33 mL (E) 42 mL

    47. What is the mass of 1 L of a gas at STP whose molar mass is 254 g/mol?

    (A) 11.3 g (B) 25.4 g (C) 30.6 g (D) 76.5 g (E) 254 g

    48. A type of ion found in aluminum oxide is:

    () X+ () X2+ (C) X3+ (D) XO32-

    () XO42-

    49. A type of ion found in potassium phosphate is:

    () X+ () X2+ (C) X3+ (D) XO32-

    () XO42-

    50. A type of ion found in sodium acetate is:

    () X+ () X2+ (C) X3+ (D) XO32-

    () XO42-

    51. How many moles of hydrogen sulfide are contained in a 35.0 g sample of this

    gas?

    () 1.03 mol () 2.06 mol (C) 6.18 mol (D) 9.45 mol () 11.3 mol

    52. What is the molar mass of ethanol (C2H5OH)?

    () 34.2 () 38.9 (C) 46.1 (D) 45.1 () 62.1

    53. Ammonia can be produced by the reaction of nitrogen and hydrogen gas.

    Suppose the reaction is carried out starting with 14 g of nitrogen and 15 g of

    hydrogen. How many grams of ammonia can be produced?

    () 17.04 g () 34.08 g (C) 51.1 g (D) 85.2 g () 102 g

    54. How many atoms of hydrogen are present in 12.0 g of water?

    () 1.1 x 1023 () 2.0 x 1023 (C) 4.0 x 1023 (D) 8.0 x 1023 () 4.8 x 1024

    55. Which compound contains the highest percent by mass of hydrogen?

    () HCI () 2 (C) 34 (D) H2SO4 () HF

    56. hydrocarbon (a compound consisting solely of carbon and hydrogen) is found

    to be 96% carbon by mass. What is the empirical formula for this compound?

    () C2H () CH2 (C) C3H (D) CH3 (E) C4H

    57. An unknown compound contains the elements C, , and O. It is known to

    contain 48% C and 4.0% by mass. The molar mass of this compound has been

    determined in the lab to have a value of 200. The molecular formula for this

    compound is:

  • SAT CHEMISTRY Dr. D. Bampilis 29

    () C2H3O2 () C4H6O 4 (C) C4H4O 3 (D) C8 H3O 6 () C8 8 O6

    58. When 7.0 g of ethene (C2H4) burns in oxygen to give carbon dioxide and water,

    how many grams of CO2 are formed?

    C2H4 + 3O2 2CO2 + 2H2O

    () 9.0 g () 22 g (C) 44 g (D) 82 g () 180 g

    59. Consider the reaction below. What mass of CF4 is formed by the reaction of 8.00

    g of methane with an excess of fluorine?

    CH4(g) + 4F2(g) CF4(g) + 4HF(g)

    () 19 g () 22 g (C) 38 g (D) 44 g () 88 g

    60. According to the reaction represented by the unbalanced equation:

    SO2(g) +O2(g) SO3(g)

    how many moles of SO2(g) are required to react completely with 1 mole of O2(g)?

    () 0.5 mol () 1 mol (C) 2 mol (D) 3 mol () 4 mol

    61. The combustion of propane, C3H8(g) , proceeds according to the equation:

    C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l)

    How many grams of water will be formed in the complete combustion of 44.0 grams of

    propane?

    () 4.50 g () 18.0 g (C) 44.0 g (D) 72.0 g () 176 g

    62. The number of oxygen atoms in 0.50 mole of KHSO4 is:

    (A) 1.2 x 1023 (B) 2.4 x 1023 (C) 3.0 x 1023 (D) 1.2 x 1024 () 2.4 x 1024

    63. Lithium reacts with water to produce hydrogen gas and lithium hydroxide. What volume

    of hydrogen collected over water at 22oC and 750 mmHg pressure is produced by the

    reaction of 0.208 g of Li? [VPH2O = 19.8 mmHg]

    (A) 367 mL (B) 378 mL (C) 735 mL (D) 755 mL

    64. Analysis by mass of a certain compound shows that it contains 14 % hydrogen and 86 %

    carbon. Which of the following is the most informative statement that can properly be made

    about the compound on the basis of these data?

    (A) It is a hydrocarbon. (B) Its empirical formula is CH2.

    (C) Its molecular formula is C2H4. (D) Its molar mass is 28 g/mol.

    (E) It contains a triple bond.

    65. In a balanced equation, the number of moles of each substance is equal

    BECAUSE

    Once the limiting reagent has been consumed, the reaction can no longer

    continue

  • SAT CHEMISTRY Dr. D. Bampilis 30

    6. Solids, Liquids and Phase Changes

    Phase Changes In order for a substance to move between the states of matter; for example, to turn from a solid into a liquid, which is called fusion, or from a gas to a liquid (vaporization), energy must be gained or lost. The heat of fusion (symbolized Hfus) of a substance is the amount of energy that must be put into the substance for it to melt. For example, the heat of fusion of water is 6.01 kJ/mol, or in other terms, 80 cal/g. The heat of vaporization, not surprisingly, is the amount of energy needed to cause the transition from liquid to gas, and it is symbolized Hvap. You will not be required to memorize heat of fusion or vaporization values for the exam. Changes in the states of matter are often shown on phase diagrams, and you will probably see at least one of two different types of phase diagrams on the SAT II Chemistry exam. Lets start with the phase diagram for water. The phase diagram for water is a graph of pressure versus temperature. Each of the lines on the graph represents an equilibrium position, at which the substance is present in two states at once. For example, anywhere along the line that separates ice and water, melting and freezing are occurring simultaneously.

    The intersection of all three lines is known as the triple point (represented by a dot and a T on the figure). At this point, all three phases of matter are in equilibrium with each other. Point X represents the critical point, and at the critical point and beyond, the substance is forever in the vapor phase. This diagram allows us to explain strange phenomena, such as why water boils at a lower temperature at higher altitudes, for example. At higher altitudes, the air pressure is lower, and this means that water can reach the boiling point at a lower temperature. Interestingly enough, water would boil at room temperature if the pressure was low enough! One final note: If we put a liquid into a closed container, the evaporation of the liquid will cause an initial increase in the total pressure of the system, and then the pressure of the system will become a constant. The value of this final pressure is unique to each liquid and is known as the liquids vapor pressure. Water has a relatively low vapor pressure because it takes a lot of energy to break the hydrogen bonds so that molecules enter the gas phase. Water and other liquids that have low vapor pressures are said to be nonvolatile. Substances like rubbing alcohol and gasoline, which have relatively high vapor pressures, are said to be volatile.

  • SAT CHEMISTRY Dr. D. Bampilis 31

    Example What happens to water when the pressure remains constant at 1 atm but the temperature changes from -10C to 75C? Explanation Looking at the phase change diagram for water and following the dashed line at 1 atm, you can see that water would begin as a solid (ice) and melt at 0C. All of the water would be in liquid form by the time the temperature reached 75C. The second type of phase change graph you might see on the SAT II Chemistry exam is called a heating curve. This is a graph of the change in temperature of a substance as energy is added in the form of heat. The pressure of the system is assumed to be held constant, at normal pressure (1 atm). As you can see from the graph below, at normal pressure water freezes at 0C and boils at 100C.

    The plateaus on this diagram represent the points where water is being converted from one phase to another; at these stages the temperature remains constant since all the heat energy added is being used to break the attractions between the water molecules. Specific Heat On the SAT II Chemistry test, you might see a diagram that looks something like this one, and you might come across a question that asks you to calculate the amount of energy needed to take a particular substance through a phase change. This would be one of the most difficult questions on the exam, but you might see something like it, or at least part of it. If you were asked to do this, you would need to use the following equation: energy (in calories) = mCp DT where m = the mass of the substance (in grams) Cp = the specific heat of the substance (in cal/g.C) DT=the change in temperature of the substance (in either Kelvin or C, but make sure all your units are compatible!) As you can see, this requires that you know the specific heat of the substance. A substances specific heat refers to the heat required to raise the temperature of 1 g of a substance by 1C. You will not be required to remember any specific heat values for the exam. Work through the example below to get a feel for how to use this equation. Example If you had a 10.0 g piece of ice at -10C, under constant pressure of 1 atm, how much energy would be needed to melt this ice and raise the temperature to 25.0C? Explanation First, the temperature of the ice would need to be raised from -10C to 0C. This would require the following calculation. The specific heat for ice is 0.485 cal/g.C. Substituting in the formula energy = mCp DT; energy = (10.0 g) (0.485 cal/g.C) (10.0C) = 48.5 cal So 48.5 calories are needed to raise temperature.

  • SAT CHEMISTRY Dr. D. Bampilis 32

    Next, we must calculate the heat of fusion of this ice: we must determine how much energy is needed to completely melt the 10 g of it. energy = mHfus energy = (10.0 g) (80 cal/g) = 800 cal So 800 cal of energy are needed to completely melt this sample of ice. Next, we need to see how much energy would be needed to raise the temperature of water from 0C to 25C. The specific heat for liquid water is 1.00 cal/g.C. So again use energy = mCp T to get energy = (10.0 g) (1.00 cal/g.C) (25.0C) = 250 cal Finally, add together all of the energies to get the total: 48.5 + 800 + 250 = about 1100 calories are needed to convert the ice to water at these given temperatures.

    You have to know:

    Liquids

    Intermolecular interaction

    Brownian movement

    Viscosity

    Surface tension

    Capillary action

    Boiling point - Vapor pressure

    Critical temperature, critical pressure

    Solids

    Crystalline - amorphous

    Sublimation - deposition

    Melting point

    Phase diagram - Triple point

    Heating curve

    Solubility - saturated solution

    Solid - gas (P, T)

    Rate of solubility (Pulverizing, Stirring, Heating)

    Solutions: Dilute - Concentration - Saturated - Unsaturated - Supersaturated

    Expressions of concentration

    % w/w

    % w/v

    % v/v

    mol solute

    Molarity(M)Volume solution (L)

    mol solute

    molality(m)1 kg solvent

    Dilution C1V1 = C2V2

    Exercises

    For questions 1 - 4:

    (A) Boyles law (B) Charles law (C) Avogadros law (D) Ideal gas law (E) Daltons law

    1. The total pressure of a gaseous mixture is equal to the sum of the partial pressures is:

  • SAT CHEMISTRY Dr. D. Bampilis 33

    2. Volume is inversely proportional to pressure is:

    3. Volume is directly proportional to temperature is:

    4. All gases have the same number of moles in the same volume at constant T and P is:

    For questions 5 -7:

    (A) Sublimation (B) Condensation (C) Evaporation (D) Deposition (E)melting

    5. Gas solid is called:

    6. Gas liquid is called:

    7. Solid gas is called:

    For questions 8 -10:

    (A) AB (B) BC (C) CD (D) DE (E) EF

    8. Which shows melting?

    9. Which shows increasing the kinetic energy of a liquid?

    10. Which shows boiling?

    For questions 11 - 19: Statement I BECAUSE Statement II

    11. The ideal gas law does not hold under low temperatures and high pressure

    BECAUSE

    Interactions between particles cannot be neglected under these conditions

    12. CO2 is able to sublimate at atmospheric pressure

    BECAUSE

    Its liquid form is impossible to produce

    13. When an ideal gas is cooled its volume will increase

    BECAUSE

    Temperature and volume are directly proportional

    14. According to the KMT, collisions between gas particles and the walls of the container are

    elastic

    BECAUSE

  • SAT CHEMISTRY Dr. D. Bampilis 34

    Gas molecules are considered volume-less particles, with no intermolecular forces, in

    constant random motion

    15. As ice absorbs heat and begins to melt, its temperature remains constant

    BECAUSE

    Changes of state bring about changes in a substances potential energy, not in its kinetic

    energy

    16. Water boils at a lower temperature at high altitudes compared to low altitudes

    BECAUSE

    The vapor pressure of water is lower at higher altitude

    17. Decreasing the volume of a system decreases pressure

    BECAUSE

    Pressure and volume are inversely related

    18. At constant pressure, a certain amount of gas will double in volume as the temperature

    is halved

    BECAUSE

    Temperature and volume are inversely proportional

    19. The volume of a gas at 100oC and 600 mmHg will be lower at STP

    BECAUSE

    Decreasing temperature and increasing pressure will cause the volume of a gas to decrease

    20. What volume would 16 g of molecular oxygen gas occupy at STP?

    (A) 5.6 L (B) 11.2 L (C) 22.4 L (D) 33.6 L (E) 44.8 L

    21. Which of the following is responsible for the abnormally high boiling point of water?

    (A) Covalent bonding (B) Hydrogen bonding (C) High polarity

    (D) Large dielectric constant (E) Low molecular weight

    22. Which of the following is (are) the weakest attractive forces?

    (A) Van der Waals (B) Coordinate covalent bonding (C) Covalent bonding

    (D) Polar covalent bonding (E) Ionic bonding

    23. What is the volume at STP of 10 L of gas initially at 546 K, 2 atm?

    (A) 5 L (B) 10 L (C) 15 L (D) 20 L (E) 25 L

    24. If one mole of H2 is compressed from 10 L to 7.5 L at constant temperature, what

    happens to the gas pressure?

    (A) It increases by 25% (B) It decreases by 25% (C) It increases by 33%

    (D) It increases by 50% (E) None of the above

  • SAT CHEMISTRY Dr. D. Bampilis 35

    25. An ideal gas in a closed inflexible container has a pressure of 6 atm and a temperature of

    27 deg C. What will be the new pressure at -73 deg C?

    (A) 2 atm (B) 3 atm (C) 4 atm (D) 8 atm (E) 9 atm

    For the next few questions, refer to the diagram below, regarding substance Z.

    26. Substance Z is at 1 atm and 200 K. If the pressure on substance Z is steadily increased

    and its temperature is kept constant, what phase change will eventually occur?

    (A) condensation (B) freezing (C) melting (D) sublimation (E) vaporization

    27. The normal boiling point of substance Z is approximately:

    (A) 100 K (B) 200 K (C) 300 K (D) 400 K (E) 500 K

    28. In what pressure range will the compound sublime?

    (A) Less than 0.5 atm (B) Between 0.5 and 1.0 (C) Between 1.0 and 2.0

    (D) Between 0.5 and 2.0 (E) This compound wont sublime

    29. Crossing line bd is:

    (A) condensation (B) melting (C) evaporation (D) sublimation (E) boiling

    30. Five liters of gas at STP have a mass of 12.5 g. What is the molecular mass of the gas?

    (A) 12.5 g/mol (B) 25.0 g/mol (C) 47.5 g/mol (D) 56.0 g/mol (E) 125 g/mol

    31. Equal molar quantities of hydrogen gas and oxygen gas are present in a closed container

    at a constant pressure. Which of the following quantities will be the same for the two gases?

    (A) Partial pressure

    (B) Partial pressure & average KE

    (C) Partial pressure & average molecular velocity

    (D) Average KE & average molecular velocity

    (E) Partial pressure, average KE, average molecular velocity

    For the next few questions: A closed 5.0 L vessel contains a sample of neon. The

    temperature inside the container is 25oC and the pressure is 1.5 atm.

    32. Which of the following expressions is equal to the moles of gas in the sample?

    (A) (1.5 x 5.0) / (0.08 x 25) (B) (0.08 x 250 / (1.5 x 5.0) (C) (1.5 x 25) / (0.08 x 5.0)

    (D) (0.08 x 298) / (1.5 x 5.0) (E) (1.5 x 5.0) / (0.08 x 298)

    33. If the neon gas in the vessel is replaced with an equal molar quantity of helium gas,

    which will be changed?

    (A) pressure (B) temperature (C) density (D) pressure & temperature

  • SAT CHEMISTRY Dr. D. Bampilis 36

    (E) temperature and density

    34. The volume was changed while temperature held constant until the pressure was 1.6

    atm. Which is equal to the new volume?

    (A) 5.0 x 1.5 / 1.6 (4.7 L) (B) 5.0 x 1.6 / 1.5 (C) 25 x 1.5 / 1.6

    (D) 0.08 x 1.6 / 1.5 (E) 0.08 x 1.5 / 1.6

    35. A flask contains three times as many moles of H2 as it does O2. If hydrogen and oxygen

    are the only gases present, what is the total pressure in the flask if the partial pressure of

    oxygen is P?

    (A) 4P (B) 3P (C) 4/3P (D) 3/4P (E) 7P

    36. The gas in a large cylinder is at a pressure of 3040 torr. Assuming constant temperature

    and ideal gas behavior, what volume of this gas could you compress into a 100 L box at 8

    atm?

    (A) 20 L (B) 200 L (C) 5000 L (D) 50,000 L (E) 500,000 L

    37. Which of the following generalizations CANNOT be made about the phase change of a

    pure substance from solid to liquid?

    (A) It involves a change in potential energy (B) It involves no change in temperature

    (C) It involves a change in kinetic energy (D) It involves a change in entropy

    (E) It may occur at different temperatures for different compounds

    38. If the pressure of a gas sample is doubled at constant temperature, the volume will be:

    (A) 4 x the original (B) 2 x the original (C) of the original

    (D) of the original (E) 1/8 of the original

    39. Three canisters, A, B, and C, are all at the same temperature, with volumes of 2.0, 4.0,

    and 6.0 L, respectively. Canister A contains 0.976 g Ar at 120 torr, Canister B contains 1.37 g

    N2 at 120 torr, and Canister C is completely empty at the start. Assuming ideality, what

    would be the pressure in canister C if the contents of A and B are completely transferred to

    C?

    (A) 180 torr (B) 330 torr (C) 675 torr (D) 0.25 atm (E) none of the above

    40. When a fixed amount of gas has its Kelvin temperature and pressure doubled, the new

    volume of the gas is:

    (A) four times greater than its original volume (B) twice its original volume

    (C) unchanged (D) one half its original volume (E) one fourth its original volume

    41. A 600 mL container holds 2 mol O2, 3 mol H2, and 1 mol He. The total pressure within the

    container is 760 torr. What is the partial pressure of O2?

    (A) 127 torr (B) 253 torr (C) 380 torr (D) 507 torr (E) 760 torr

    42. An ideal gas has a volume of 10 L at 20 deg C and 750 mmHg. Which of the following

    expressions is needed to determine the volume of the same amount of gas at STP?

  • SAT CHEMISTRY Dr. D. Bampilis 37

    (A) 10 x (750/760) x (0/20) (B) 10 x (750/760) x (293/273) (C) 10 x (760/750) x (0/20)

    (D) 10 x (760/750) x (273/293) (E) 10 x (750/760) x (273/293)

    43. What volume does a sample of 1.50 x 1023 atoms of helium at STP represent?

    (A) 5.6 L (B) 11.2 L (C) 17.8 L (D) 22.4 L (E) none of the above

    44. Which of the following will always decrease the volume of a gas?

    i. Decrease the pressure with the temperature held constant

    ii. Increase the pressure with a temperature decrease

    iii. Increase the temperature with a pressure increase

    (A) I only (B) II only (C) I and III (D) II and III only (E) I, II and III

    45. A gas has a volume of 10 L at 50 deg C and 200 mmHg. What conversion factor is needed

    to give a volume at STP?

    (A) 10 x (0/50) x (200/760) (B) 10 x (0/50) x 760/200)

    (C) 10 x (273/323) x (200/760) (D) 10 x (273/323) x (760/200)

    (E) 10 x (323/273) x (760/200)

    46. The temperature above which a liquid cannot exist is indicated by:

    (A) the triple point (B) the critical point (C) the eutectic point (D) the boiling point

    (E) the sublimation point

    47. A change of phase never accompanies:

    (A) a change in volume (B) a change in pressure (C) a change in temperature

    (D) a change in density (E) a change in structure

    48. The relationship P1V1 = P2V2 is:

    (A) Boyles law (B) Charless law (C) Van der Waals law

    (D) the combined gas law (E) the ideal gas law

    49. The rate of diffusion of hydrogen gas as compared to that of oxygen gas is:

    (A) as fast (B) identical (C) twice as fast (D) four times as fast

    (E) eight times as fast

    50. The ratio of the rate of diffusion of oxygen to hydrogen is:

    (A) 1:2 (B) 1:4 (C) 1:8 (D) 1:16 (E) 1:32

    51. Standard conditions using a Kelvin thermometer are:

    (A) 760 torr, 273 K (B) 760 torr, 273 K, 1 L (C) 760 torr, 0 K

    (D) 0 torr, 0 K (E) 0 torr, 273 K, 1 L

    52. The relation between the pressure and the volume of a gas at constant temperature is

    given by:

    (A) Boyles law (B) Charless law (C) the combined gas law

    (D) the ideal gas law (E) none of the above

  • SAT CHEMISTRY Dr. D. Bampilis 38

    53. The relation between the absolute temperature and volume of a gas at constant

    pressure is given by:

    (A) Boyles law (B) Charless law (C) the combined gas law

    (D) the ideal gas law (E) none of the above

    54. The relation between the pressure, volume and absolute temperature is given by:

    (A) Boyles law (B) Charless law (C) the combined gas law

    (D) the ideal gas law (E) none of the above

    55. At a certain temperature and pressure, ice, water and steam are found to coexist at

    equilibrium. This pressure and temperature corresponds to:

    (A) the critical temperature (B) the critical pressure (C) the sublimation point

    (D) the triple point (E) two of the above

    56. How many atoms are present in 22.4 L of O2 at STP?

    (A) 3 x 1023 (B) 6 x 1023 (C) 9 x 1023 (D) 12 x 1023 (E) 15 x 1023

    57. gas at STP that contains 6.02 x 1023 atoms and forms diatomic molecules will occupy:

    (A) 11.2 L (B) 22.4 L (C) 33.6 L (D) 67.2 L (E) 1.06 quarts

    58. Inelastic collisions occur in:

    (A) Real and ideal gases (B) Ideal gases and fusion reactions

    (C) Real gases and fusion reactions (D) Real gases (E) Ideal gases

    59. The extremely high melting point of diamond (carbon) may be explained by the:

    (A) network covalent bonds (B) ionic bonds (C) hydrogen bonds

    (D) van der Waals forces (E) none of the above

    60. The phase transition from gas to solid is called:

    (A) condensation (B) evaporation (C) polymerization (D) sublimation

    61. Which aqueous solution freezes at the lowest temperature?

    (A) 0.30 m C2H5OH (B) 0.25 m KNO3 (C) 0.20 m CaBr2 (D) 0.10 m FeCl3

    62. The value of which concentration unit for a solution changes with temperature?

    (A) molarity (B) molality (C) mole fraction (D)mass percentage

    63. What is the molality of a solution made by dissolving 36.0 g of glucose (C6H12O6, M =

    180.2) in 64.0 g of H2O?

    (A) 0.0533 (B) 0.200 (C) 0.360 (D) 3.12

    64. When a nonvolatile solute is dissolved in a volatile solvent, which characteristic is greater

    for the solution than for the solvent?

    (A) boiling point (B) freezing point (C) rate of evaporation (D) vapor pressure

  • SAT CHEMISTRY Dr. D. Bampilis 39

    65. Which aqueous solution has the highest osmotic pressure at 25oC? (Assume all ionic

    compounds ionize completely in solution.)

    (A) 0.1 M Al2(SO4)3 (B) 0.1 M Na2CO3 (C) 0.2 M KMnO4 (D) 0.3 M C6H12O6

    66. A student prepares four 0.10 M solutions, each containing one of the solutes below.

    Which solution has the lowest freezing point?

    (A) CaCl2 (B) KOH (C) NaC2H3O2 (D) NH4NO3

    67. Ethanol (C2H5OH, M = 46) and methanol (CH3OH, M = 32) form an ideal solution when

    mixed. What is the vapor pressure of a solution prepared by mixing equal masses of ethanol

    and methanol? (The vapor pressures of ethanol and methanol are 44.5 mmHg and 88.7

    mmHg, respectively.)

    (A) 133 mmHg (B) 70.6 mmHg (C) 66.6 mmHg (D) 44.5 mmHg

    68. A sample of H2 collected over H2O at 23C and a pressure of 732 mmHg has a volume of

    245 mL. What volume would the dry H2 occupy at 0C and 1 atm pressure?

    [vp H2O at 23C = 21 mmHg]

    (A) 211 mL (B) 218 mL (C) 224 mL (D) 249 mL

    69. An ice cube at an unknown temperature is added to 25.0 g of liquid H2O at 40.0C. The

    final temperature of the 29.3 g equilibrated mixture is 21.5C. What was the original

    temperature of the ice cube? [Cp (J/g.oC) water = 4.184, ice = 2.06, Hfusion = 333 J/g]

    (A) -6.5C (B) -13.1C (C) -35.3C (D) -56.8C

    70. A sample of a volatile liquid is introduced to an evacuated container with a movable

    piston. Which change occurs as the piston is raised? (Assume some liquid remains.)

    I. The fraction of the molecules in the gas phase increases.

    II. The pressure in the container decreases.

    (A) I only (B) II only (C) Both I and II (D) Neither I nor II

    71. Which point on the phase diagram represents the normal boiling point?

    (A) point A (B) point B (C) point C (D) point D

  • SAT CHEMISTRY Dr. D. Bampilis 40

    72. According to the following information, in what physical state(s) does bromine exist at -

    7.4oC and 400 mmHg?

    [Triple point -7.3oC, 44 mmHg, Liquid density 3.1 g.cm3, Solid density 3.4 g.cm3]

    (A) solid only (B) liquid only (C) liquid and solid only (D) gas, liquid, and solid

    73. The molarity of a solution that is composed of 80.00 g of sodium hydroxide

    dissolved in 2.0 L of solution is:

    () 1.0 () 2.0 (C) 4.0 (D) 40.0 () 160.0

    74. The molarity of a solution obtained when 50.0 mL of 6.0 HCl is diluted to a

    final volume of 300.0 mL is:

    () 0.01 () 0.10 (C) 0.20 (D) 0.30 () 1.0

    75. In the laboratory, a sample of hydrogen is collected by water displacement. The

    sample of hydrogen has a volume of 25 mL at 24.0oC and a barometric pressure for

    the day of 758 mmHg. What is the pressure of the dry gas at this temperature? (The

    vapor pressure of water at 24.0oC is 22.4 mmHg.)

    () 455 mmHg () 470 mmHg (C) 736 mmHg (D) 758 mmHg () 780 mmHg

    76. At the molecular level, the factor that determines whether a substance will be a solid,

    liquid, or gas is the balance between the:

    (A) kinetic energies of the molecules and their intermolecular forces.

    (B) potential energies of the molecules and their intermolecular forces.

    (C) kinetic energies of the molecules and their intramolecular forces.

    (D) potential energies of the molecules and their intramolecular forces.

    77. The critical temperature of carbon dioxide is 304.3 K. Which statement is true about the

    behavior of carbon dioxide above this temperature?

    (A) Solid, liquid and gaseous carbon dioxide are in equilibrium above this temperature.

    (B) Liquid and gaseous carbon dioxide are in equilibrium above this temperature.

    (C) Liquid carbon dioxide does not exist above this temperature.

    (D) Carbon dioxide molecules do not exist above this temperature.

    78. The solubility of KClO3 at several temperatures is shown in the accompanying diagram.

  • SAT CHEMISTRY Dr. D. Bampilis 41

    A student mixes 10.0 g of KClO3 with 45.0 g of H2O and stirs it for a long time at 60C until

    the solution is completely clear then allows it to cool slowly to 20C where it remains clear.

    Which statement about the final clear mixture at 20C is correct?

    (A) It is a saturated solution.

    (B) It is an unsaturated solution and can be made saturated by decreasing the temperature.

    (C) It is an unsaturated solution and can be made saturated by increasing the temperature.

    (D) It is a supersaturated solution.

    79. Which of the following solutions would probably have the highest boiling

    point?

    () 0.100 m () 0.100 m Na2SO4 (C) 0.100 m C6H12O6

    (D) 0.200 m CaCl2 () 0.200 m CH3CH2OH

    80. To determine whether a water solution of Na2S2O3 at room temperature is

    supersaturated, one can:

    (A) heat the solution to its boiling point.

    (B) add water to the solution.

    (C) add a crystal of Na2S2O3 to the solution.

    (D) acidify the solution.

    (E) cool the solution to its freezing point.

    81. Which of the following must be measured in order to calculate the molality of a

    solution?

    . Mass of the solute. . Mass of the solvent.

    III. Total volume of the solution.

    (A) I only (B) I and III only (C) II and III only (D) I and II only

    () , , and III

    82. Which of the fo l lowing so lutes an d solvents would be expected to form

    stable solutions?

    SOLUTE SOLVENT

    I . Ethanol Water

    I I . Salt Water

    I I I . Oi l Vinegar

    IV. Oi l Gasol ine

    ( ) only (B) I and I I only (C) I I I only (D) I , I I , and I I I only

    (E) I , I I , and IV only

    83. Using the sketch of the phase diagram for water given below, determine which

    of the following statements is incorrect:

  • SAT CHEMISTRY Dr. D. Bampilis 42

    () The triple point is point . This is the point at which all three phases are in

    equilibrium with one another.

    () The line AB is the line representing the solid-liquid equilibrium line. Anywhere

    along this line the substance could melt or freeze.

    (C) The slope of line AB is negative. This slope indicates that the solid is much

    denser than the liquid.

    (D) Line AD represents the phase changes of sublimation and deposition.

    () Line AC represents where the substance would condense and vaporize.

    84. A thermometer is placed in a test tube containing a melted pure substance. As slow

    cooling occurs, the thermometer is read at regular intervals until well after the sample has

    solidified. Which of the following types of graphs is obtained by plotting temperature versus

    time for this experiment?

    For questions 85 - 87: Statement I BECAUSE Statement II

    85. The solubility of carbon dioxide in a soft drink decreases with a decrease in pressure.

    BECAUSE

    The solubility of a gas generally increases with an increase in temperature.

    86. Most ionic solids have high melting points.

    BECAUSE

    Ionic solids are made up of positive and negative ions held together by

    electrostatic attractions.

  • SAT CHEMISTRY Dr. D. Bampilis 43

    87. The rate at which sugar dissolves in water increases with stirring.

    BECAUSE

    Stirring exposes the surface of a solute crystal to a less concentrated layer of solution.

    88. Which statement about the triple point of a substance is correct?

    (A) The triple point for a substance varies with the pressure.

    (B) The three phases (solid, liquid, gas) have the same density.

    (C) The three phases (solid, liquid, gas) are in equilibrium.

    (D) The three phases (solid, liquid, gas) are indistinguishable in appearance.

    89. Under certain conditions CO2 melts rather than sublimes. To which transition in the

    phase diagram does this change correspond?

    (A) A B (B) A C (C) B C (D) C B

  • SAT CHEMISTRY Dr. D. Bampilis 44

    7. Reaction Types

    Net Ionic Equations Net ionic equations are equations that show only the soluble, strong electrolytes reacting (these are represented as ions) and omit the spectator ions, which go through the reaction unchanged. When you encounter net ionic equations on the SAT II Chemistry test, youll need to remember the following solubility rules, so memorize them! Also keep in mind that net ionic equations, which are the bare bones of the chemical reaction, usually take place in aqueous environments. Here are those solubility rules:

    1. Most alkali metal compounds and NH4+ compounds are soluble.

    2. Cl-, Br-, I- compounds are soluble, except when they contain Ag+, Hg22+, or Pb2+.

    3. F- compounds are soluble, except when they contain group 2A metals. 4. NO3

    -, ClO3-, ClO4

    - and CH3COO- compounds are soluble.

    5. SO42- compounds are soluble, except when they include Ca2+, Sr2+, Ba2+, Ag+, Pb2+, or Hg2

    2+. 6. CO3

    2-, PO43-, C2O4

    2-, CrO42-, S2-, OH-, and O2- compounds are insoluble.

    7. Group 2A metal oxides are classified as strong bases even though they are not very soluble. Here are some additional rules about common reaction types that you should be familiar with for the exam:

    If an insoluble precipitate or gas can be formed in a reaction, it probably will be. Oxides (except group 1A) are insoluble, and when reacted with water, they form either acids

    (nonmetal oxides) or bases (metal oxides). There are six strong acids that completely ionize: HCl, HBr, HI, HNO3, H2SO4, HClO4. All other

    acids are weak and are written together, as molecules. The strong bases that ionize are oxides and hydroxides of group 1A and 2A metals. All other

    oxides and hydroxides are considered weak and written together, as molecules. Now try writing some net ionic equations, using the rules above. Example Write the net ionic equation for a mixture of solutions of silver nitrate and lithium bromide. Explanation

    Ag+ + NO3- + Li+ + Br-

    This is a double replacement reaction. Both compounds are soluble, so everything ionizes. If anything is formed, it will come from recombining the inside two ions with the outside two ions to make LiNO3 and AgBr. If either of them is insoluble, a precipitate will be formed, and the ions that react to form it will be in our net ionic equation; the other ions are spectators and should be omitted! As we said, the two possible products are lithium nitrate and silver bromide. Since halides are soluble except those containing silver, mercury, or lead, we have a precipitate of silver bromide, and our net ionic equation looks like this:

    Ag+ + Br- AgBr Example Hydrochloric acid and sodium hydroxide are mixed. Write the net ionic equation. Explanation This is a mixture of a strong acid and a strong base, so each ionizes completely.

    H+ + Cl- + Na+ + OH- The two possible compounds formed are sodium chloride, which is soluble, and water, which is molecular; thus water is the only product in our net ionic equation.

    H+ + OH- H2O Example Chlorine gas is bubbled into a solution of potassium iodide; write the net ionic equation.

  • SAT CHEMISTRY Dr. D. Bampilis 45

    Explanation This one is a single replacement, so you need to consider the activity series. Since halogens are involved, you can determine their activity by using the periodic table: Cl is more active than I.

    Cl2 + K+ + I-

    Remember that halogen is diatomic and that all potassium compounds are soluble. The resulting compound is also soluble, so K+ is a spectator and is left out of the final equation.

    Cl2 + I- I2 + Cl

    -

    You have to know:

    Combination - Synthesis *

    Zn(s) + S(s) ZnS(s)

    2Na(s) + Cl2(g) 2NaCl(s)

    2Mg(s) + O2(g) 2MgO(s)

    H2(g) + Cl2(g) 2HCl(g)

    H2(g) + 1/2O2(g) H2O(l)

    3H2(g) + N2(g) 2NH3(g)

    C(s) + O2(g) CO2(g)

    4Al(s) + 3O2(g) 2Al2O3(s)

    Decomposition

    2HgO(s) 2Hg(s) + O2(g)

    2H2O(l) 2H2(g) + O2(g) (electrolysis)

    2KClO32MnO 2KCl + 3O2(g)

    Single replacement

    Li, Rb, K, Ba, Sr, Ca, Na

    2Na(s) + 2H2O(l)cold 2NaOH(aq) + H2(g)

    Ca(s) + H2SO4(l) CaSO4(s) + H2(g)

    *(Ba(s) + O2(g) BaO)

    Mg, Al, Mn, Zn, Cr, Fe, Cd

    Mg(s) + H2O(g) MgO(s) + H2(g)

    Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

    *(2Fe(s) + O2(g) 2FeO(s))

    Co, Ni, Sn, Pb

    Pb(s) + HCl(aq) PbCl2(aq) + H2(g)

    * (2Cu(s) + O2(g) 2CuO(s))

    Ag, Pt, Au

    Double replacement

    AB + CD AD + CB

    D or CB: insoluble precipitate gas weak electrolyte (H2O..)

    Soluble:

    Na+, K+, NH4+, NO3

    -, CH3COO-, HCO3

    -, ClO3-

    Insoluble:

    AgCl, HgCl, PCl2 (sol in hot water)

    BaSO4, BaSO4, PbSO4, Ag2SO4, HgSO4, SrSO4

  • SAT CHEMISTRY Dr. D. Bampilis 46

    All carbonates, posphates except K+, Na+, NH4+

    All sulfides, hydroxides except K+, Na+, NH4+, Ba2+, Ca2+

    Gases: HX, H2S, HCN, SO2(H2SO3), CO2(H2CO3), NH3(NH4OH)

    Exercises

    1. Indium reacts wit h bromine to form In Br 3 . In the ba lanced equat ion for

    this react ion, t he coeff ic ient o f the indium ( ) bromide is :

    ( ) 1 ( ) 2 (C) 3 (D) 4 ( ) 6

    2 . What is the sum of the coefficients of the following equation when it is balanced?

    2 (SO 4 ) 3 + Ca(OH) 2 AI (OH) 3 + CaSO 4

    ( ) 5 ( ) 6 (C) 7 (D) 8 ( ) 9

    3 . When the following equation is balanced, what is the sum of the coefficients?

    2 (CO 3 ) 3 + Mg(OH) 2 AI (OH) 3 + MgCO 3

    ( ) 3 ( ) 4 (C) 8 (D) 9 ( ) 10

    4. When the equation for the reaction: BCl3(g) + H2(g) HCl(g) + B(s) is balanced and all

    coefficients are reduced to lowest whole-number terms, the coefficient for HCl is:

    () 1 () 2 (C) 3 (D) 4 () 6

    5. The balanced net ionic equation for the reaction of aluminum sulfate and sodium

    hydroxide contains which of the following terms?

    ( ) 3 3 + ( a q ) ( ) OH-

    ( a q ) (C) 3OH-(aq) (D) 2

    3+(aq) () 2(OH)3(s)

    6. When solutions of phosphoric acid and iron () nitrate react, which of the

    following terms will be present in the balanced molecular equation?

    () 3(aq) () 3HNO3(aq) (C) 2FePO4(s) (D) 3FePO4(s) () 2HNO3(aq)

    7. Which solid does not react with a small amount of 3 M HNO3?

    (A) calcium carbonate (B) manganese (II) sulfide (C) potassium sulfite (D) silver chloride

    8. The reaction of silver ion with chloride ion in water solution is a(n):

    (A) Precipitation (B) Oxidation-reduction (C) Distillation

    (D) Hydration (E) Condensation

    9. The reaction of iron filings with powdered sulfur is a(n):

    (A) Precipitation (B) Oxidation-reduction (C) Distillation

    (D) Hydration (E) Condensation

    10. A student prepares a 100 mL aqueous solution containing a small amount of (NH4)2SO4 and a second 100 mL solution containing a small amount of Nal, then mixes the two

    solutions. Which statement describes what happens?

    (A) Both compounds dissolve and remain in solution when the two solutions are mixed.

    (B) Both compounds dissolve initially but NH4I precipitates when the solutions are mixed.

  • SAT CHEMISTRY Dr. D. Bampilis 47

    (C) Both compounds dissolve initially but Na2SO4 precipitates when the solutions are mixed.

    (D) The NaI dissolves but the (NH4)2SO4 does not. There is no change upon mixing.

    11. Mixing which pair of 0.10 M solutions produces two precipitates that cannot be

    separated from one another by filtration?

    (A) aluminum chloride and copper(II) nitrate

    (B) strontium bromide and lead(II) acetate

    (C) magnesium perchlorate and lithium carbonate

    (D) barium hydroxide and copper(II) sulfate

    12. A colored gas is observed with which combination?

    (A) calcium hydride and water (B) lead metal and nitric acid

    (C) sodium carbonate and sulfuric acid (D) zinc sulfide and hydrochloric acid

    13. Mixing which combination produces a gaseous product?

    (A) solid ammonium nitrate and solid calcium hydroxide

    (B) copper metal and 0.10 M hydrochloric acid

    (C) solutions of barium hydroxide and 0.10 M sulfuric acid

    (D) solutions of aluminum nitrate and sodium chloride

  • SAT CHEMISTRY Dr. D. Bampilis 48

    8. Thermodynamics

    You have to know:

    Endothermic reaction:

    C(s) + O2(g) CO2(g), 0

    or C(s) + O2(g) CO2(g) + q

    Exothermic reacrion:

    N2(g) + O2(g) 2NO(g), 0

    or N2(g) + O2(g) 2NO(g) - q

    Standard enthalpy of formation o

    f

    Standard enthalpy of combustion o

    c

    Calorimetry: q = mc q = mL (phase change)

    Hesss law o o

    rxn f f(products) (reac tan ts)

    o o

    rxn c c(reac tan ts) (products)

    Bond dissociation energy (bond length)

    Bond broken: Energy absorbed

    Bond formed: Energy evolved o

    rxn broken formedD D Entropy S > 0

    solid liquid gas

    T

    mixing different particles

    (gas)

    Rxnn

    Exercises

    For questions 1- 3: Refer to the following potential energy diagram & the choices below:

    -26.4 kcal -67.6 kcal C(s) CO(g) CO2(g) (A) -94.0 kcal (B) -26.4 kcal (C) -67.6 kcal (D) C(s) + O2(g) () CO2(g)

    1. What is the H of the reaction to form CO from C + O2?

    G = - S

    - - + Spontaneous

    - - spontaneous at low T

    + + - Nonspontaneous

    + + spontaneous at high T

  • SAT CHEMISTRY Dr. D. Bampilis 49

    2. What is the H of the reaction to form CO2 from CO + O2?

    3. What is the H of the reaction to form CO2 from C + O2?

    For questions 4 - 6: Refer to the heating curve:

    4. In which part of the curve is the state only a solid?

    5. In which part is the heat to change the state greatest?

    6. In which part is the heat to change the temperature greatest?

    For questions 7 - 8:

    7. Which letter shows the potential energy of the products?

    8. Which letter shows the enthalpy change (H) of the reaction?

    For questions 9 - 13: Refer to the heating curve for H2O below:

    9. Where is the temperature of H2O changing at 1

    oC/g.cal?

    10. Which region indicates a solid?

    11. Which region indicates a liquid?

    12. Which region indicates a gas?

    13. Which region indicates a liquid and a gas?

  • SAT CHEMISTRY Dr. D. Bampilis 50

    For questions 14 - 20: Statement I BECAUSE Statement II

    14. An exothermic reaction has a positive H

    BECAUSE

    Heat is released in an exothermic reaction

    15. A calorimeter can be used to measure the amount of heat lost or absorbed in a process

    BECAUSE

    Calorimeters can be used to measure heat lost or gained by a system and its surroundings

    16. The freezing of water is an exothermic process

    BECAUSE

    Energy is released when covalent bonds are formed

    17. An increase in entropy leads to a decrease in randomness

    BECAUSE

    The low energy state of ordered crystals has high entropy

    18. An exothermic reaction has a positive H value

    BECAUSE

    Heat must be added to an exothermic reaction for the reaction to occur

    19. Covalent bonds must be broken for a liquid to boil

    BECAUSE

    Heat is released when a liquid changes into a gas

    20. The temperature of a substance always increases as heat energy is added to it

    BECAUSE

    The average kinetic energy of the particles in a system increases with an increase in

    temperature

    21. How much heat is given off when 8 g of hydrogen reacts in: 2H2 + O2 2H2O?

    H = -115.60 kcal

    (A) -57.8 kcal (B) -115.6 kcal (C) -173.4 kcal (D) -231.2 kcal () -462.4 kcal

    22. A reaction that absorbs heat is:

    (A) endothermic (B) an equilibrium process (C) spontaneous

    (D) non-spontaneous () exothermic

    23. The change in heat energy for a reaction is best expressed as a change in:

    (A) Enthalpy (H) (B) Absolute temperature (T) (C) Specific heat (c)

    (D) Entropy (S) () Kinetic energy (KE)

    24. When 1 mole of sulfur burns to form SO2, 1300 calories are released. When 1 mole of

    sulfur burns to form SO3, 3600 calories are released. What is H when 1 mole of SO2 burns

    to form SO3?

  • SAT CHEMISTRY Dr. D. Bampilis 51

    (A) 3900 cal (B) -1950 cal (C) 1000 cal (D) -500 cal () -2300 cal

    25. When the temperature of a 20 gram sample of water is increased from 10oC to 30oC, the

    heat absorbed by the water is

    (A) 600 cal (B) 30 cal (C) 400 cal (D) 20 cal () 200 cal

    26. How many g of CH4 produce 425.6 kcal in: CH4 + 2O2 CO2 + 2H2O + 212.8 kcal

    (A) 8 g (B) 16 g (C) 24 g (D) 32 g () 64 g

    27. 10 g of liquid at 300 K are heated to 350 K. The liquid absorbs 6 kcal. What is the specific

    heat of the liquid (in cal/g.oC)?

    (A) 6 (B) 120 (C) 12 (D) 600 () 60

    28. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + 800 kJ. If a mole of O2(g) is consumed in the

    reaction, what energy is produced?

    (A) 200 kJ (B) 400 kJ (C) 800 kJ (D) 1200 kJ () 1600 kJ

    29. What is Hrxn for the decomposition of 1 mole of NaClO3? Hf = -85.7 kcal/mol for

    NaClO3(s) and Hf = -98.2 kcal/mol for NaCl(s)

    (A) -183.9 kcal (B) -91.9 kcal (C) +45.3 kcal (D) +22.5 kcal () -12.5 kcal

    30. What is the heat of combustion of one mole of C2H4?

    Compound Hf (kcal/mol) H2O(g) -57.8 C2H4(g) 12.5 CO2(g) -94.1

    (A) +316.3 kcal (B) -12.5 kcal (C) -291.3 kcal (D) -316.3 kcal () -57.8 kcal

    31. Given 2Na(s) + Cl2(g) 2NaCl(s) + 822 kJ, how much heat is released if 0.5 mol of

    sodium reacts completely with chlorine?

    (A) 205.5 kJ () 411 kJ (C) 822 kJ (D) 1644 kJ () 3288 kJ

    32. Calculate the approximate amount of heat necessary to raise the temperature of 50.0

    grams of liquid water from 10.0oC to 30.0o