Chemistry H1: Form TR9.6A
Name ______________________________
REVIEW
Date _________________ Period _____
Test Review # 9
Hybrid Orbitals. The native orbitals found in an atom in the free state cannot
always account for the geometry of the compounds formed from the atom.
σ bonds
Mixing of native atomic orbitals to allow bonding to occur is known as
hybridization. [1] When there are four effective pairs of electrons, one s
orbital combines with three p orbitals to form four equivalent sp3 hybrid
orbitals. sp3 hybrid orbitals are tetrahedral. [2] Whenever an atom is
sp2
sp2
surrounded by three effective electron pairs to form a trigonal planar
molecule, a set of sp2 hybrid orbitals is required.. Combination of one s orbital
sp2
and two p orbitals to form an sp2 hybrid gives the appropriate 120Eangle. In
2
σ bonds
forming the sp orbital, one p orbital is not used, and is oriented perpendicular
to the plane of the sp2 orbitals each of the three sp2 orbitals forms bonds by
sharing a pair of electrons in an area centered on a line between the two
atoms. These are called sigma (σ) bonds. The double bond is formed in the
p orbitals
space above and below the σ bond by the p orbital perpendicular to the sp2
orbitals. This is called a pi (π) bond. [3] sp hybridization enables two effective
pairs of electrons to bond at 180E. One s and one p are hybridized to form two
sp hybrid orbitals at a 180E angle. Two p orbitals remain. The hybrid orbitals
form sigma bonds and the p orbitals form pi bonds. [4] dsp3 hybridization
enables a trigonal bipyrimidal arrangement for five pairs of electrons
surrounding a central atom, while d2sp3 hybridization enables an octahedral
pi bond
arrangement for six pairs of electrons surrounding a central atom.
Metallic Bonding. Metals have low ionization energies. This means they hold
onto electrons loosely. As a result, in a metal crystal, the valence electrons
move easily and do not belong to any single atom. Since the atoms in the
crystal do not
hold on to their
own valence
sigma bond
electrons, they
become like
cations in a sea
of mobile
electrons. The attraction between the cations and the electrons holds the metal
crystal together. Because of this, metals are lustrous, flexible, good conductors
of heat and electricity, and are solids at room temperature except for mercury.
Polar Molecules. Electronegativity differences between 0.4 and 1.7 are found
in molecules with polar bonds. These molecules can be polar depending on
their shapes. Molecules with polar bonds distributed symmetrically are nonpolar. Asymmetrical molecules with polar bonds are polar.
Water is polar. An imaginary line can be drawn through a water molecule separating the positive pole from the negative pole. This is
because the charges are distributed asymmetrically. Carbon dioxide is nonpolar because the electronegative oxygens are distributed
symmetrically around the carbon. (O=C=O)
Types of Bonds. Pure substances can be held together by ionic bonds, covalent bonds, metallic bonds, or intermolecular forces. All ionic
substances are crystalline solids. Diamonds are also crystalline solids, but they are made of pure carbon. Large crystals such as diamond
or sand (SiO2) that have a network of covalent bonds are called macromolecules or network solids. Smaller compounds containing
covalent bonds are called molecules. The molecules of a substance may be attracted to each other to form solids or liquids by
intermolecular forces. These are often called molecular compounds. Molecular solids are softer than covalent solids (network solids) and
ionic solids, because intermolecular forces are weaker than chemical bonds. If the substance is polar, it is held together by dipole-dipole
attractions. If the polar substance contains hydrogen atoms attached to either oxygen, nitrogen, or fluorine atoms, it forms especially strong
dipole-dipole attractions called a hydrogen bonds. Hydrogen bonds are responsible for the three dimensional shapes of many proteins
because the large protein molecule folds in such a way that hydrogens in one part of the molecule are close to oxygens or nitrogens in
another part of the molecule. Nonpolar molecules are attracted to each other only by the weakest intermolecular forces called Van der
Waal’s forces or London Dispersion Forces. London dispersion forces increase with molecular size. Scientists believe they are caused
by temporary dipoles that form as the electrons shift positions.
C
C
C
C
Chemistry H1: Form TR9.6A
REVIEW
Test Review 9
Page 2
PHASES OF MATTER
Comparing Solids, Liquids, and Gases. Solids are substances with a definite shape and volume. The particles of solids vibrate about fixed
positions, held in place by large forces of attraction. Liquids have a definite volume, but their shape is determined by their container. The
particles of a liquid roll and slide over each other. Both the shape and volume of a gas are determined by the container. This is because
the particles move independently, and spread out to fill the container. Gases are mostly empty space, and they can be compressed.
Kinetic Molecular Theory. Matter is composed of particles that are in constant motion (kinetic energy). There are forces of attraction
between particles that depend on the distance between the particles. The further apart the particles are, the smaller the forces of attraction
between them are. The higher the temperature (average kinetic energy) is, the faster the particles move. The Kinetic Molecular Theory
explains the phases. In solids the forces of attraction between particles are larger than in other phases. As a result, the particles are held
relatively close together in fixed positions, vibrating back and forth. Therefor the shape and volume are not determined by the container.
In liquids the forces of attraction between particles are moderate compared to other phases. The particles can move from place to place
but cannot separate from each other and move independently, so they roll and slide over each other. The particles are pulled downhill by
gravity causing the liquid to seek its own level, so the shape is determined by the container but the volume is not. In gases the forces of
attraction between particles are weaker than in other phases. The particles can move from place to place independently of each other
because they do NOT attract or repel each other. The particles are relatively far apart. The volume of the particles is small compared to
the space between them. Gases tend to spread out to fill their container. Therefor both the shape and volume are determined by the
container.
PHASE CHANGES
Heating a substance in a given phase causes the temperature to increase. Increasing the temperature
causes particles to move faster and collide harder. This causes the particles to rebound harder moving
them further apart. Larger distances between particles weakens the forces of attraction between them.
When the forces of attraction are weak enough, the distance between the particles increases markedly
and the phase changes. As a result, a solid melts, and a liquid evaporates. The reverse happens when
a substance cools, so a gas condenses, and a liquid freezes.
Phase Diagram. A phase diagram shows phases of a substance in a closed system as a function of
temperature and pressure. The points on the phase diagram are [a] the triple point where all three
phases of mater coexist; [b] the critical temperature above which vapor cannot be liquefied no matter
what pressure is applied; [c] the critical pressure, the pressure required to form a liquid at the critical
temperature; and [d] the critical point defined by the critical temperature and pressure.
Heating/Cooling Curve. The temperature does not change during a phase change. The heat energy
absorbed or lost does not result in a change in kinetic energy. Instead, there is a change in potential
energy due to the change in distance between the particles. The
Freezing/Melting point is the temperature at which the solid and liquid
Phase diagram for water
phase exist in equilibrium. The heat of fusion is the amount of heat
needed to change a unit mass of a substance from a solid to a liquid at
a constant temperature and 1 atm of pressure. For water it is 333.6 J/g.
The boiling point is the temperature at which the vapor pressure is
equal to the surrounding pressure. The heat of vaporization is the
amount of heat needed to change a unit mass of a substance from a
liquid to a gas at a constant temperature and 1 atm of pressure. For
water it is 2259 J/g.
Vapor Pressure. When water evaporates, it changes from a liquid to
a gas called water vapor. Water vapor takes up more space than an
equal mass of liquid water. As a result, in a closed container, the vapor
that forms can exert a significant amount of pressure. This pressure is
known as vapor pressure. Even in an open container, the vapor is
confined by the air pressing down on it. Some of it collects at the
surface and exerts pressure. Occasional high energy molecules at the
water’s surface escape. That is why the water eventually evaporates.
But for a water to expand and form vapor bubbles throughout the liquid as it does when it boils, the vapor has to exert as much pressure
as the blanket of air confining it. As a liquid is heated, more of it turns into vapor, and the vapor pressure increases. When the vapor
pressure reaches atmospheric pressure, the liquid boils. Under greater external pressure, the liquid boils at a higher temperature.
Test Review 9
Chemistry H1: Form TR9.6A
REVIEW
Page 3
Answer the questions below by circling the number of the correct response
1. Which of the following molecules contains an sp2–sp2 sigma bond?
(1) CH4 (2) C2H2 (3) C2H4 (4) C2H6 (5) none of these
14. Which substance exists as a metallic crystals? (1) Ar (2) SiO2
(3) Au (4) CO2
2. What type of hybridization is found in CH2O? (1) sp3 (2) sp2 (3) sp
(4) dsp3 (5) d2sp3
15. Mobile electrons are a distinguishing characteristic of (1) an ionic
bond (2) a metallic bond (3) an electrovalent bond (4) a covalent
bond
3. What type of hybridization explains the shape of an octahedral
molecule? (1) sp3 (2) sp2 (3) sp (4) dsp3 (5) d2sp3
4. Which of the following molecules has sp hybrid bonds? (1) CH4
(2) C2H2 (3) C2H4 (4) C2H6 (5) none of these
5. Consider the following Lewis structure.
H
H
O
H
H
C
C
C
C
3
2
1
H
H
6. Which molecule is nonpolar and has a symmetrical shape? (1) HCI
(2) CH4 (3) H2O (4) NH3
7. Which formula represents a polar molecule? (1) CH4 (2) Cl2 (3) NH3
(4) N2
8. Which structural formula represents a nonpolar symmetrical
molecule?
H
H
(2) H
O
H
C
H
H
17. Which of the following is an example of hydrogen bonding?
(1) H2(R) (2) I2(s) (3) CH3OH(R) (4) C8H18(R)
18. The boiling point increases as you go down the halogen family
because of the increase in (1) van der Waal’s forces, (2) metallic
properties, (3) polarity, (4) covalent bonding.
Which statement about the molecule is false? (1) There are 10
sigma and 2 pi bonds. (2) C–2 is sp2 hybridized with bond angles
of 120E. (3) Oxygen is sp3 hybridized. (4) This molecule contains
28 valence electrons. (5) There are some H–C–H bond angles of
about 109E in the molecule.
(1)
16. Which element consists of positive ions immersed in a "sea" of
mobile electrons.? (1) sulfur (2) calcium (3) nitrogen (4) chlorine
(3)
H
F
(4)
H
20. Which 5.0-milliliter sample of NH3 will take the shape of and
completely fill a closed 100.0-milliliter container? (1) NH3(s)
(2) NH3(R) (3) NH3(g) (4) NH3(aq)
21. Which of the following has the strongest forces of attraction?
(1) CO2(s) (2) CO2(R) (3) CO2(g) (4) CO2(aq)
22. Which of the following can be compressed under pressure?
(1) I2(s) (2) I2(R) (3) I2(g) (4) I2(aq)
23. Which 1.5-liter sample of salt does NOT take the shape of its
container? (1) NaCl(s) (2) NaCl(R) (3) NaCl(g) (4) NaCl(aq)
N
H
19. In the family of compounds including H2O, H2S, H2Se, and H2Te,
water has the highest boiling point because it has the greatest
(1) van der Waal’s forces, (2) metallic bonding, (3) polarity,
(4) covalent bonding.
H
9. Why is NH3 classified as a polar molecule? (1) It is a gas at STP.
(2) H—H bonds are nonpolar. (3) Nitrogen and hydrogen are both
nonmetals. (4) NH3 molecules have asymmetrical charge
distributions.
10. Which statement best explains why carbon tetrachloride (CCl4) is
nonpolar? (1) Each carbon-chloride bond is polar. (2) Carbon and
chlorine are both nonmetals. (3) Carbon tetrachloride is an organic
compound. (4) The carbon tetrachloride molecule is symmetrical.
24. A 25.0 mL sample of water is poured from a 50.0 mL graduated
cylinder to a 100.0 mL graduated cylinder. The volume of the water
(1) increases, (2) decreases, (3) remains the same.
25. As ice melts at standard pressure, its temperature remains at 0EC
until it has completely melted. Its potential energy (1) decreases
(2) increases (3) remains the same
26. When water freezes, each gram loses an amount of heat equal to
its heat of (1) fusion (2) vaporization (3) sublimation (4) reaction
11. Which substance will conduct electricity in both the solid phase and
the liquid phase? (1) AgCl (2) H2 (3) Ag (4)HCl
27. As the temperature of a liquid increases, its vapor pressure
(1) decreases (2) increases (3) remains the same
12. What type of bonds are present in a strip of magnesium ribbon?
(1) covalent (2) metallic (3) ionic (40 van der Waals
28. Which change of phase represents fusion? (1) gas to liquid
(2) gas to solid (3) solid to liquid (4) liquid to gas
13. Which substance, in the solid state, is the best conductor of
electricity? (1) Ag (2) NaCl (3) I2 (4) CO2
29. Which substance readily sublimes at room temperature?
(1) H2O(R) (2) O2(g) (3) Fe(s) (4) CO2(s)
Test Review 9
Chemistry H1: Form TR9.6A
REVIEW
30. Which change of phase represents sublimation?
(1) H2O(g) ÿ H2O(R) (2) H2O(R) ÿ H2O(s) (3) CO2(s) ÿ CO2(g)
(4) CO2(s) ÿ CO2(R)
Page 4
Base your answers to questions 40 and 41 on the diagram below which
represents a substance being from a solid to a gas, the pressure
remaining constant
31. Which change of phase is exothermic? (1) gas to liquid (2) solid to
liquid (3) solid to gas (4) liquid to gas
32. The heat of fusion for ice is 333.6 joules per gram. Adding 333.6
joules of heat to one gram of ice at STP will cause the ice to
(1) increase in temperature (2) decrease in temperature (3) change
to water at a higher temperature (4) change to water at the same
temperature
33. Which term represents the change of a substance from the solid
phase to the liquid phase? (1) condensation (2) vaporization
(3) evaporation (4) fusion
44. Which of the following has the highest boiling point at a pressure of
40 kPa? (1) propanone (2) ethanoic acid (3) ethanol (4) water
3
2
3
3
2
3
1
24.
25.
26.
27.
28.
29.
30.
Answers
3
2
1
2
3
4
3
31.
32.
33.
34.
35.
36.
37.
1
4
4
4
4
2
1
38.
39.
40.
41.
42.
43.
44.
3
2
4
3
2
4
2
39. The temperature at which a substance can exist as a solid, liquid, and
gas simultaneously is the (1) melting point, (2) triple point, (3) boiling
point, (4) critical point.
43. At what temperature will water boil at a pressure of 10 kPa?
(1) 25EC (2) 10EC (3) 101.3EC (4) 45EC
17.
18.
19.
20.
21.
22.
23.
38. A substance sublimes at standard temperature and pressure. What
could be done to cause the substance to melt? (1) increase the
temperature (2) decrease the temperature (3) increase the pressure
(4) decrease the pressure.
42. At what pressure will ethanol boil at 90EC? (1) 75 kPa (2) 150 kPa
(3) 101.3 kPa (4) 200 kPa
4
4
3
2
1
3
2
2
37. The heat of fusion of a substance is the energy measured during a
(1) phase change (2) temperature change (3) chemical change
(4) pressure change
Answer questions 42-44 by referring to Table H on the reference tables.
9.
10.
11.
12.
13.
14.
15.
16.
36. Which sample contains particles arranged in regular geometric
pattern? (1) CO2(R) (2) CO2(s) (3) CO2(g) (4) CO2(aq)
41. Between points B and C the substance exists in (1) the solid state,
only (2) the liquid state, only (3) both the solid and liquid states
(4) neither the solid nor the liquid state
3
2
5
2
3
2
3
2
35. The energy required to change a unit mass of a solid to a liquid at
constant temperature is called its heat of
(1) formation
(2) vaporization (3) combustion (4) fusion
40. The substance begins to boil at point (1) E (2) B (3) C (4) D
1.
2.
3.
4.
5.
6.
7.
8.
34. When the vapor pressure of a liquid in an open container equals the
atmospheric pressure, the liquid will (1) freeze (2) crystallize
(3) melt (4) boil