Midterm Review Packet With Questions
Midterm Review Packet With Questions
Midterm Review Packet With Questions
Midterm
Review Packet
Name:
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Table of Contents
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Topic 1: Matter, Its Properties and Changes Outline
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5. Properties can be physical or chemical. Physical properties describe
those characteristics that can be observed with the senses or
measured. Chemical properties describe how the substance interacts
with other substances.
Distinguish between chemical and physical properties.
One of the more useful properties is density. The density equation is on
Table T; D=m/V.
Some common properties of the elements are found on Table S, such
as melting and boiling points.
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Matter – Cut from Jan 2007 – Jan 2008 Exams
1. A sample composed only of atoms having the 8. Which element is a solid at STP and a good
same atomic number is classified as conductor of electricity?
(1) a compound (3) an element (1) iodine (3) nickel
(2) a solution (4) an isomer (2) mercury (4) sulfur
2. A dilute, aqueous potassium nitrate solution 9. The table below shows mass and volume data
is best classified as a for four samples of substances at 298 K and
(1) homogeneous compound 1 atmosphere.
(2) homogeneous mixture
(3) heterogeneous compound
(4) heterogeneous mixture
4. Which statement describes a chemical property of Which two samples could consist of the same
hydrogen gas? substance?
(1) Hydrogen gas burns in air. (1) A and B (3) B and C
(2) Hydrogen gas is colorless. (2) A and C (4) C and D
(3) Hydrogen gas has a density of 0.000 09 g/cm3
at STP. 10. Bronze contains 90 to 95 percent copper and 5 to
(4) Hydrogen gas has a boiling point of 20. K at 10 percent tin. Because these percentages can vary,
standard pressure. bronze is classified as
(1) a compound (3) a mixture
5. Which element has the greatest density at STP? (2) an element (4) a substance
(1) calcium (3) chlorine
(2) carbon (4) copper 11. At STP, which list of elements contains a solid, a
liquid, and a gas?
6. Which statement describes a chemical property (1) Hf, Hg, He (3) Ba, Br2, B
of the element magnesium? (2) Cr, Cl2, C (4) Se, Sn, Sr
(1) Magnesium is malleable.
(2) Magnesium conducts electricity. 12. A 10.0-gram sample of which element has the
(3) Magnesium reacts with an acid. smallest volume at STP?
(4) Magnesium has a high boiling point. (1) aluminum (3) titanium
(2) magnesium (4) zinc
7. Matter that is composed of two or more
different elements chemically combined in a fixed 13. At room temperature, a mixture of sand and
proportion is classified as water can be separated by
(1) a compound (3) a mixture (1) ionization (3) filtration
(2) an isotope (4) a solution (2) combustion (4) sublimation
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14. Which particle diagram represents a sample of one compound, only?
15. A 1.00-mole sample of neon gas occupies a volume of 24.4 liters at 298 K and 101.3 kilopascals. Calculate the
density of this sample. Your response must include both a correct numerical setup and the calculated result.
In an investigation, a dripless wax candle is massed and then lighted. As the candle burns, a small amount of
liquid wax forms near the flame. After 10 minutes, the candle’s flame is extinguished and the candle is allowed to
cool. The cooled candle is massed.
16. Identify one physical change that takes place in this investigation.
17. State one observation that indicates a chemical change has occurred in this investigation.
•
18. Draw a particle diagram showing the change from solid wax to liquid wax. Use “ ” for particles of wax.
Draw separate diagrams for the liquid and the solid states.
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Base your answers to questions 19 through 21 on the particle diagrams below, which show atoms and/ or
molecules in three different samples of matter at STP.
20. When two atoms of y react with one atom of z, a compound forms. Using the number of atoms shown
in sample 2, what is the maximum number of molecules of this compound that can be formed?
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Topic 2: Atomic Concepts Outline
1. The modern model of the atom has evolved over a long period of time through the
work of many scientists.
Dalton’s Model:
Elements are made of atoms
Atoms of an element are the same.
Compounds are formed from combinations of atoms.
Rutherford Experiment
Bombarded gold foil with alpha particles. Showed atoms
were mostly empty space with small, dense positively
charged nucleus.
Bohr Model
Small, dense, positively charged nucleus surrounded by electrons in circular orbits.
Wave-Mechanical Model (Modern Atomic Theory)
Small, dense, nucleus positively charged nucleus
surrounded by electrons moving in “electron cloud”.
“Orbitals” are areas where an electron with a certain amount of energy is most likely
to be found.
2. Each atom is made of a positively charged nucleus with one or more orbiting,
negatively charged electrons.
6. The mass of a proton is 1 amu. The mass of a neutron is 1 amu. The mass of an
electron is almost 0 amu.
The mass of an atom is contained in its nucleus.
Theatomic massof an atom is equal to the total number of protons and neutrons.
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8. When the electron gains a specific amount of energy, it moves to a higher orbital
and is in the “excited state”.
You can recognize an excited state electron configuration. If the configuration does
not match that on the Periodic Table for that number of electrons, then it is an
excited state.
9. When an electron returns from a higher energy state to a lower energy state,
it emits a specific amount of energy usually in the form of light. This can be used
to identify an element (bright line spectrum).
The instrument used to see the bright line spectrum is called a spectroscope.
10. The outermost electrons are called valence electrons. These affect the chemical
properties of the element.
Atoms with a filled valence level are stable (noble gases).
Most elements can have up to 8 electrons in their valence level. The exceptions are H and
He, which can have only 2 valence electrons.
Atoms form bonds in order to fill their valence levels.
You can use Lewis structures to show the configuration of the valence electrons.
11. Atoms of the same element all contain the same number of protons.
Changing the number of protons changes the atom into a different
element.The atomic number is the number of protons in an atom of an
element.
12. Isotopes are atoms with equal numbers of protons but different numbers of
neutrons.
Isotopes of an element have the same atomic number (protons only), but different
atomic masses (protons + neutrons).
13. The average atomic mass of an element is the weighted average of its naturally
occurring isotopes.
You need to know how to do the calculation of “weighted atomic mass” given isotope
masses and percent abundances.
14. When an atom gains an electron, it becomes a negative ion and its radius
increases.
15. When an atom loses an electron, it becomes a positive ion and its radius
decreases.
1. Experiments performed to reveal the structure of atoms led scientists to conclude that an
atom’s
(1) positive charge is evenly distributed throughout its volume
(2) negative charge is mainly concentrated in its nucleus
(3) mass is evenly distributed throughout its volume
(4) volume is mainly unoccupied
3. An experiment in which alpha particles were used to bombard thin sheets of gold foil led
to the conclusion that an atom is composed mostly of
(1) empty space and has a small, negatively charged nucleus
(2) empty space and has a small, positively charged nucleus
(3) a large, dense, positively charged nucleus
(4) a large, dense, negatively charged nucleus
4. What is the atomic number of an element that has six protons and eight neutrons?
(1) 6 (2) 2 (3) 8 (4) 14
5. An atom of fluorine has a mass of 19 atomic mass units. The total number of protons
and neutrons in its nucleus is
(1) 9 (2) 10 (3) 19 (4) 28
6. What is the total number of protons contained in the nucleus of a carbon-14 atom?
(1) 6 (2) 8 (3) 12 (4) 14
8. Which of these elements has an atom with the most stable outer electron configuration?
(1) Ne (2) Cl (3) Ca (4) Na
9. How many electrons are in the outermost principal energy level of an atom of carbon in
the ground state?
(1) 6 (2) 2 (3) 3 (4) 4
11. What is the electron configuration of a sulfur atom in the ground state?
(1) 2-4 (2) 2-6 (3) 2-8-4 (4) 2-8-6
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12. The nucleus of which atom contains 48 neutrons?
(1) 32 S (2) 48Ti (3) 85 Rb (4) 112Cd
16 22 37 48
14. When an atom loses an electron, the atom becomes an ion that is
(1) positively charged and gains a small amount of mass
(2) positively charged and loses a small amount of mass
(3) negatively charged and gains a small amount of mass
(4) negatively charged and loses a small amount of mass
15. In which pair of elements do the nuclei of the atoms contain the same number of
neutrons?
(1) 3
7
Li and 4
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Be (3) 1123 Na and 1224 Mg
(2) 14
7N and 16
8O (4) 1632 S and 1735Cl
16. The characteristic spectral lines of elements are caused when electrons in an excited
atom move from
(1) lower to higher energy levels, releasing energy
(2) lower to higher energy levels, absorbing energy
(3) higher to lower energy levels, releasing energy
(4) higher to lower energy levels, absorbing energy
17. Which Lewis electron-dot structure is drawn correctly for the atom it represents?
18. When a lithium atom forms a Li+ ion, the lithium atom
(1) gains a proton (3) loses an electron
(2) loses a proton (4) gains an electron
19. What is the total number of electrons in the valence shell of an atom of aluminum in the
ground state?
(1) 8 (2) 2 (3) 3 (4) 10
20. An electron in an atom moves from the ground state to an excited state when the energy
of the electron
(1) increases (2) decreases (3) remains the same
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21. During a flame test, ions of a specific metal are heated in the flame of a gas burner. A
characteristic color of light is emitted by these ions in the flame when the electrons
(1) emit energy as they move to higher energy levels
(2) emit energy as they return to lower energy levels
(3) gain energy as they move to higher energy levels
(4) gain energy as they return to lower energy levels
Base your answers to questions 23 and 24 on the information and the bright-line spectra
represented below.
Many advertising signs depend on the production of light emissions from gas-filled glass
tubes that are subjected to a high-voltage source. When light emissions are passed
through a spectroscope, bright-line spectra are produced.
24. Explain the production of an emission spectrum in terms of the energy states of an electron.
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Atomic Concepts Review – Cut from Jan 2007 – Jan 2008 Exams
1. Which subatomic particles are located in the 6. Which isotopic notation identifies a metalloid
nucleus of a neon atom? that is matched with the corresponding number
(1) electrons and positrons of protons in each of its atoms?
(2) electrons and neutrons
q(3) protons and neutrons
(4) protons and electrons
of ?
(1) 26 (3) 57
(2) 31 (4) 83
14. What was concluded about the structure of the 19. Which two notations represent different
atom as the result of the gold foil experiment? isotopes of the same element?
(1) A positively charged nucleus is
surrounded by positively charged particles.
(2) A positively charged nucleus is
surrounded by mostly empty space.
(3) A negatively charged nucleus is surrounded
by positively charged particles.
(4) A negatively charged nucleus is surrounded
by mostly empty space.
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Base your answers to questions 20 through 22 on the information below.
The accepted values for the atomic mass and percent natural abundance of each naturally occurring
isotope of silicon are given in the data table below.
21. Show a correct numerical setup for calculating the atomic mass of Si. [1]
22. A scientist calculated the percent natural abundance of Si-30 in a sample to be 3.29%.
Determine the percent error for this value. [1]
23. Write one electron configuration for an atom of silicon in an excited state.
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Base your answers to questions 24 through 26 on the information below.
24. Identify one piece of information shown in the electron-shell diagrams that is not
shown in the Lewis electron-dot diagrams. [1]
25. Determine the mass number of the magnesium atom represented by the electron-shell
diagram. [1]
26. Explain why Lewis electron-dot diagrams are generally more suitable than electron-shell diagrams for
illustrating chemical bonding. [1]
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Topic 3: Periodic Table Outline
4. The atomic mass is the sum of protons and neutrons in the nucleus.
The mass number given on the periodic table is a weighted average of the different
isotopesof that element.
Electrons do not significantly add to the atomic mass.
6. Elements can be classified by their properties and their location on the Periodic
Table as metals, non-metals, metalloids, and noble gases.
10. Elements of the same period have the same number of occupied energy levels.
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11. Elements of the same group have the same valence configuration and similar
chemical properties.
Group 1 elements other than H arealkali
metals.Group 2 elements arealkali earth
metals.
Group 17 elements arehalogens.
Alkali metals, alkali earth metals, and halogens all are highly reactive and do not exist
asfree elements in nature (they are all found in compounds).
Group 18 elements arenoble or inert gases. These elements have filled valence levels
andare do not normally react with other substances.
14. Some elements may exist in two or more forms in the same phase. These forms
differ in their molecular or crystal structure, hence their different properties. These
different forms are called “allotropes,”
Ex: Solid carbon exists in three different forms: graphite, diamond (a network solid) and
coal.
Ex: the element oxygen can exist in two different forms: O2 gas and ozone (O3 gas)
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Periodic Table – Practice Questions
2. Elements in a given period of the Periodic Table contain the same number of
(1) protons in the nucleus (3) electrons in the outermost level
(2) neutrons in the nucleus (4) occupied principal energy levels
5. Which two elements have chemical properties that are most similar?
(1) Cl and Ar (3) K and Ca
(2) Li and Na (4) C and N
6. Which of the following Period 4 elements has the most metallic characteristics?
(1) Ca (2) Ge (3) As (4) Br
7. If M represents an alkali metal, what is the formula for the compound formed by M and
oxygen?
(1) MO2 (2) M2O (3) M2O3 (4) M3O2
8. As the elements in Group 15 are considered in order of increasing atomic number, which
sequence in properties occurs?
(1) nonmetal, metalloid, metal (3) metal, metalloid, nonmetal
(2) metalloid, metal, nonmetal (4) metal, nonmetal, metalloid
10. As elements of Group 15 of the Periodic Table are considered in order from top to
bottom, the metallic character of the atoms of each successive element generally
(1) increases (2) decreases (3) remains the same
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11. Which statement best describes Group 2 elements as they are considered in order from
top to bottom of the Periodic Table?
(1) The number of principal energy levels increases, and the number of valence
electrons increases.
(2) The number of principal energy levels increases, and the number of valence
electrons remains the same.
(3) The number of principal energy levels remains the same, and the number
of valence electrons increases.
(4) The number of principal energy levels remains the same, and the number
of valence electrons decreases.
13. Which Group 16 element when combined with hydrogen forms a compound that would
exhibit the strongest hydrogen bonding?
(1) selenium (3) oxygen
(2) tellurium (4) sulfur
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Periodic Table – Cut from Jan 2007 – Jan 2008 Exams
1. Which element is a solid at STP and a good 8. Which list of elements consists of metalloids,
conductor of electricity? only?
(1) iodine (3) nickel (1) B, Al, Ga (3) O, S, Se
(2) mercury (4) sulfur (2) C, N, P (4) Si, Ge, As
2. Which element has both metallic and nonmetallic 9. Which general trend is found in Period 2 on the
properties? Periodic Table as the elements are considered in
(1) Rb (3) Si order of increasing atomic number?
(2) Rn (4) Sr (1) decreasing atomic mass
(2) decreasing electronegativity
3. The carbon atoms in graphite and the carbon (3) increasing atomic radius
atoms in diamond have different (4) increasing first ionization energy
(1) atomic numbers
(2) atomic masses
(3) electronegativities 10. Which two characteristics are associated
(4) structural arrangements with metals?
(1) low first ionization energy and low
4. Atoms of which element have the greatest electronegativity
tendency to gain electrons? (2) low first ionization energy and high
(1) bromine (3) fluorine electronegativity
(2) chlorine (4) iodine (3) high first ionization energy and low
electronegativity
5. Which statement describes a chemical property (4) high first ionization energy and high
of the element magnesium? electronegativity
(1) Magnesium is malleable.
(2) Magnesium conducts electricity.
(3) Magnesium reacts with an acid. 11. Which element is most chemically similar to
(4) Magnesium has a high boiling point. chlorine?
(1) Ar (3) Fr
6. Which statement explains why sulfur is (2) F (4) S
classified as a Group 16 element?
(1) A sulfur atom has 6 valence electrons. 12. Which grouping of circles, when considered in
(2) A sulfur atom has 16 neutrons. order from the top to the bottom, best represents the
(3) Sulfur is a yellow solid at STP. relative size of the atoms of Li, Na, K, and Rb,
(4) Sulfur reacts with most metals. respectively?
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14. An atom of argon rarely bonds to an atom of 17. Solid samples of the element phosphorus can
another element because an argon atom has be white, black, or red in color. The variations in
(1) 8 valence electrons color are due to different
(2) 2 electrons in the first shell (1) atomic masses
(3) 3 electron shells (2) molecular structures
(4) 22 neutrons (3) ionization energies
(4) nuclear charges
15. The elements on the Periodic Table
are arranged in order of increasing 18. Lithium and potassium have similar chemical
(1) boiling point (3) atomic number properties because the atoms of both elements
(2) electronegativity (4) atomic mass have the same
(1) mass number
(2) atomic number
16. Which element is classified as a nonmetal? (3) number of electron shells
(1) Be (3) Si (4) number of valence electrons
(2) Al (4) Cl
Elements with atomic numbers 112 and 114 have been produced and their IUPAC names are pending approval.
However, an element that would be put between these two elements on the Periodic Table has not yet been
produced. If produced, this element will be identified by the symbol Uut until an IUPAC name is approved.
21. Determine the charge of an Uut nucleus. Your response must include both the numerical value and the sign of
the charge. [1]
22. Identify one element that would be chemically similar to Uut. [1]
22
Base your answers to questions 23 through 26 on the information below, which describes the proposed discovery of
element 118.
In 1999, a nuclear chemist and his team announced they had discovered a new element by crashing
krypton atoms into lead. The new element, number 118, was assigned the name ununoctium and the symbol Uuo.
One possible isotope of ununoctium could have been Uuo-291.
However, the discovery of Uuo was not confirmed because other scientists could not reproduce the
experimental results published by the nuclear chemist and his team. In 2006, another team of scientists claimed that
they produced Uuo. This claim has yet to be confirmed.
Adapted from Discover January 2002
23. Based on atomic number, in which group on the Periodic Table would element 118 be placed? [1]
24. What would be the total number of neutrons present in a theoretical atom of Uuo-291?[1]
25. What would be the total number of electrons present in a theoretical atom of Uuo-291? [1]
26. Explain why being able to reproduce scientific results is an important component of scientific research. [ 1]
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Base your answers to questions 27 through 30 on the information below.
The table below lists physical and chemical properties of six elements at standard pressure that
correspond to known elements on the Periodic Table. The elements are identified by the code letters, D, E, G,
J, L, and Q.
27. What is the total number of elements in the “Properties of Six Elements at Standard Pressure” table that
are solids at STP? [1]
28. An atom of element G is in the ground state. What is the total number of valence electrons in this atom?
[1]
29. Letter Z corresponds to an element on the Periodic Table other than the six listed elements. Elements
G, Q, L, and Z are in the same group on the Periodic Table, as shown in the diagram below.
Based on the trend in the melting points for elements G, Q, and L listed in the
“Properties of Six Elements at Standard Pressure” table, estimate the melting point of element Z, in degrees
Celsius. [1]
30. Identify, by code letter, the element that is a noble gas in the “Properties of Six Elements at Standard
Pressure” table. [1]
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Topic 4: Formulas & Names, Equations, Moles,
Molar Mass, & Types of Reactions Outline
7. The molar mass of a substance is the sum of the atomic masses of its atoms. The
molar mass (gram formula mass) equals the mass of one mole of that substance.
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Equations & Stoichiometry – Practice Questions
3. What is the total number of moles of atoms present in 1 gram formula mass of
Pb(C2H3O2)2?
(1) 9 (2) 14 (3) 3 (4) 15
10. What is the total number of moles of sulfur atoms in 1 mole of Fe2(SO4)3?
(1) 1 (2) 15 (3) 3 (4) 17
What is the coefficient of Al2(SO4)3 when the equation is completely balanced using the
smallest whole-number coefficients?
(1) 1 (2) 2 (3) 3 (4) 4
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12. Given the unbalanced equation:
When this equation is correctly balanced using smallest whole numbers, what is the
coefficient of O2 (g)?
(1) 6 (2) 2 (3) 3 (4) 4
4 NH3 + 5 O2 4 NO + 6 H2O
What is the total number of moles of NO produced when 1.0 mole of O2 is completely
consumed?
(1) 1.0 mole (2) 1.2 moles (3) 0.80 mole (4) 4.0 moles
What is the total number of moles of HCl (g) produced when 3 moles of H2 (g) is
completely consumed?
(1) 5 moles (2) 2 moles (3) 3 moles (4) 6 moles
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Formulas, Equations & Stoichiometry Review –
Cut from Jan 2007 – Jan 2008 Exams
1. Which equation shows conservation of 8. Which formula represents an ionic
atoms? compound?
(1) H2 + O2 H2O (1) H2 (3) CH3OH
(2) H2 + O2 2 H2O (2) CH4 (4) NH4Cl
(3) 2 H2 + O2 2 H2O
(4) 2 H2 + 2 O2 2 H2O 9. What is the total number of different elements
present in NH4NO3?
2. Which substance can be broken down by a (1) 7 (3) 3
chemical change? (2) 9 (4) 4
(1) antimony (3) hexane
(2) carbon (4) sulfur 10. Which formula represents lead (II)
chromate?
3. What is the gram formula mass of Ca3(PO4)2? (1) PbCrO4 (3) Pb2CrO4
(1) 248 g/mol (3) 279 g/mol (2) Pb(CrO4)2 (4) Pb2(CrO4)3
(2) 263 g/mol (4) 310 g/mol
11. Which particle diagram represents a sample
4. In which compound is the ratio of metal ions of one compound, only?
to nonmetal ions 1 to 2?
(1) calcium bromide
(2) calcium oxide
(3) calcium phosphide
(4) calcium sulfide
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12. An atom in the ground state contains a total 15. Given the balanced equation representing
of 5 electrons, 5 protons, and 5 neutrons. a reaction:
Which Lewis electron-dot diagram 4NH3 + 5O2 4NO + 6H2O
represents this atom? What is the minimum number of moles of O2
that are needed to completely react with
16 moles of NH3?
(1) 16 mol (3) 64 mol
(2) 20. mol (4) 80. mol
13. Given the balanced equation representing 16. Element X reacts with iron to form two
the reaction between propane and oxygen: different compounds with the formulas FeX
C3H8 + 5O2 3CO2 + 4H2O and Fe2X3.
According to this equation, which ratio of To which group on the Periodic Table does
oxygen to propane is correct? element X belong?
(1) Group 8 (3) Group 13
(2) Group 2 (4) Group 16
19. A hydrated compound contains water molecules within its crystal structure. The percent composition
by mass of water in the hydrated compound CaSO 4•2H2O has an accepted value of 20.9%. A student
did an experiment and determined that the percent composition by mass of water in CaSO4•2H2O was
21.4%.
Calculate the percent error of the student’s experimental result. Your response must include both a
correct numerical setup and the calculated result. [2]
20. Write the empirical formula for the compound C8H18. [1]
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Base your answers to questions 21 through 23 on the information below.
Some dry chemicals can be used to put out forest fires. One of these chemicals is NaHCO3. When
NaHCO3(s) is heated, one of the products is CO2(g), as shown in the balanced equation below.
21. Show a correct numerical setup for calculating the percent composition by mass of carbon in the
product Na2CO3. [1]
23. Determine the total number of moles of CO2(g) produced when 7.0 moles of NaHCO3(s)
is completely reacted. [1]
moles
Rust on an automobile door contains Fe2O3(s). The balanced equation representing one of the reactions
between iron in the door of the automobile and oxygen in the atmosphere is given below.
25. Identify the type of chemical reaction represented by this equation. [1]
26. Determine the gram-formula mass of the product of this reaction. [1]
30
Base your answers to questions 28 through 30 on the information below.
Ozone gas, O3, can be used to kill adult insects in storage bins for grain without damaging the
grain. The ozone is produced from oxygen gas, O2, in portable ozone generators located near the
storage bins. The concentrations of ozone used are so low that they do not cause any
environmental damage. This use of ozone is safer and more environmentally friendly than a
method that used bromomethane, CH3Br. However, bromomethane was more effective than ozone
because CH3Br killed immature insects as well as adult insects.
Adapted From: The Sunday Gazette (Schenectady, NY) 3/9/03
28. Determine the total number of moles of CH3Br in 19 grams of CH3Br (gram-formula mass =
95 grams/mol). [1]
30. Based on the information in the passage, state one advantage of using ozone instead of bromomethane
for insect control in grain storage bins. [1]
31
Base your answers to questions 31 through 33 on the information below.
A hydrate is a compound that has water molecules within its crystal structure. The
formula for the hydrate CuSO4•5H2O(s) shows that there are five moles of water for every
one mole of CuSO4(s). When CuSO4•5H2O(s) is heated, the water within the crystals is
released, as represented by the balanced equation below.
A student first masses an empty crucible (a heat-resistant container). The student then
masses the crucible containing a sample of CuSO4•5H2O(s). The student repeatedly heats
and masses the crucible and its contents until the mass is constant. The student’s recorded
experimental data and calculations are shown below.
31. Identify the total number of significant figures recorded in the calculated mass of CuSO4•5H2O(s). [1]
32. In the space below, use the student’s data to show a correct numerical setup for calculating the percent
composition by mass of water in the hydrate. [1]
33. Explain why the sample in the crucible must be heated until the constant mass is reached. [1]
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Topic 5: Bonding Outline
8. The electronegativity difference between two bonded atoms can determine the
type of bond and its polarity.
0.0 = non-polar covalent
0.0 -1.7 = polar covalent
1.7+ = ionic
9. Bonding guidelines:
Metals react with nonmetals to form ionic compounds.
Nonmetals bond with nonmetals to form covalent compounds
(molecules).Ionic compounds with polyatomic ions have both ionic and
covalent bonds.
11. Metallic bonding occurs between atoms of metal. The valence electrons are
loosely held by all atoms in a mobile “sea” of valence electrons.
This type of bonding accounts for some of the unique properties of metals, such as
theirability to conduct electricity, luster, and malleability.
1. The forces between atoms that create chemical bonds are the result of interactions
between
(1) nuclei (3) protons and electrons
(2) electrons (4) protons and nuclei
2. According to Reference Table S, which sequence correctly places the elements in order of
increasing ionization energy?
(1) H Li Na K (3) O S Se Te
(2) I Br Cl F (4) H Be Al Ga
4. If the electronegativity difference between the elements in compound NaX is 2.0, what is
element X?
(1) bromine (2) chlorine (3) fluorine (4) oxygen
6. Which type of bond exists between an atom of carbon and an atom of fluorine?
(1) ionic (2) metallic (3) polar covalent (4) nonpolar covalent
10. The primary forces of attraction between water molecules in H2O (l) are
(1) ionic bonds (3) molecule-ion attractions
(2) hydrogen bonds (4) van der Waals forces
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11. Which structure represents a polar molecule?
(1) H – H (3)
(2) (4)
12. Which electron dot diagram represents a molecule that has a polar covalent bond?
1)
2)
3)
4)
36
Bonding Review – Cut from Jan 2007 – Jan 2008 Exams
1. Given the balanced equation: 8. When sodium and fluorine combine to
produce the compound NaF, the ions formed have
I + I → I2 the same electron configuration as atoms of
Which statement describes the process (1) argon, only
represented by this equation? (2) neon, only
(1) A bond is formed as energy is absorbed. (3) both argon and neon
(2) A bond is formed and energy is released. (4) neither argon nor neon
(3) A bond is broken as energy is absorbed.
(4) A bond is broken and energy is released. 9. Atoms of which element have the greatest
tendency to gain electrons?
2. An oxygen molecule contains a double bond (1) bromine (3) fluorine
because the two atoms of oxygen share a total of (2) chlorine (4) iodine
(1) 1 electron (3) 3 electrons
(2) 2 electrons (4) 4 electrons 10. Which polyatomic ion contains the greatest
number of oxygen atoms?
3. A double carbon-carbon bond is found in a (1) acetate (3) hydroxide
molecule of (2) carbonate (4) peroxide
(1) pentane (3) pentyne
(2) pentene (4) pentanol 11. Which formula represents an ionic compound?
(1) H2 (3) CH3OH
4. At STP, fluorine is a gas and bromine is a liquid
because, compared to fluorine, bromine has (2) CH4 (4) NH4Cl
(1) stronger covalent bonds
12. Which liquid has the highest vapor pressure at
(2) stronger intermolecular forces
75°C?
(3) weaker covalent bonds
(1) ethanoic acid (3) propanone
(4) weaker intermolecular forces
(2) ethanol (4) water
5. Which term indicates how strongly an
13. Given the balanced equation representing a
atom attracts the electrons in a chemical bond?
reaction:
(1) alkalinity
Cl2(g) → Cl(g) + Cl(g)
(2) atomic mass
What occurs during this change?
(3) electronegativity
(1) Energy is absorbed and a bond is broken.
(4) activation energy
(2) Energy is absorbed and a bond is formed.
(3) Energy is released and a bond is broken.
6. Magnesium nitrate contains chemical bonds
(4) Energy is released and a bond is formed.
that are
(1) covalent, only
(2) ionic, only 14. At standard pressure, a certain compound has
a low boiling point and is insoluble in water. At
(3) both covalent and ionic
STP, this compound most likely exists as
(4) neither covalent nor ionic
(1) ionic crystals
7. A solid substance is an excellent conductor of (2) metallic crystals
electricity. The chemical bonds in this substance (3) nonpolar molecules
are most likely (4) polar molecules
(1) ionic, because the valence electrons are
shared between atoms
(2) ionic, because the valence electrons are mobile
(3) metallic, because the valence electrons are
stationary
(4) metallic, because the valence electrons are
mobile
37
15. Which group on the Periodic Table of the
Elements contains elements that react with
oxygen to form compounds with the general
formula X2O?
(1) Group 1 (3) Group 14
(2) Group 2 (4) Group 18
38
19. Explain, in terms of electronegativity, why a P–Cl bond in a molecule of PCl5 is more polar than a P–
S bond in a molecule of P2S5. [1]
The graph below shows the relationship between boiling point and molar mass at standard pressure for
pentane, hexane, heptane, and nonane.
20. Octane has a molar mass of 114 grams per mole. According to this graph, what is the boiling point of
octane at standard pressure? [1] ________________________
21. State the relationship between molar mass and the strength of intermolecular forces for the selected
alkanes. [1]
39
Base your answers to questions 22 through 24 on the information below.
The particle diagrams below represent the reaction between two nonmetals, A2 and Q2.
22. Using the symbols A and Q, write the chemical formula of the product. [1] ________________
23. Identify the type of chemical bond between an atom of element A and an atom of element Q. [1]
24. Compare the total mass of the reactants to the total mass of the product. [ 1]
25. Explain, in terms of molecular structure or distribution of charge, why a molecule of methane is nonpolar.
[1 ]
26. A liquid boils when the vapor pressure of the liquid equals the atmospheric pressure on the surface of the
liquid. Using Reference Table H, determine the boiling point of water when the atmospheric pressure is 90.
kPa. [1]
40
Base your answers to questions 27 through 30 on the information below.
Have you ever seen an insect called a water strider “skating” across the surface of a calm pond? Have you ever
“floated” a sewing needle on the water in a glass? If you have, then you’ve observed one of water’s many
amazing properties. Water’s surface tension keeps the water strider and the sewing needle from sinking into
the water. Simply stated, the surface tension is due to the forces that hold the water molecules together.
Without these intermolecular forces, the water strider and the sewing needle would sink below the surface of
the water. The surface tension of water at various temperatures is given in the data table below.
27. On a piece of graph paper, plot the data from the data table. Circle and connect the five points. [ 1]
28. According to your graph, what is the surface tension of water at 60.°C? [1] ___________ mN/m
29. State the relationship between the surface tension and the temperature of water. [1]
30. The surface tension of liquid tetrachloromethane, CCl4, at 25°C is 26.3 millinewtons/ meter (mN/m).
Compare the intermolecular forces between molecules of CCl4 to the intermolecular forces between
molecules of water, H2O, at 25°C. [1]
41
Topic 6 Overview
2. The Law of Conservation of Energy states that energy can not be lost or
destroyed, only changed from one form to another.
3. Heat is a transfer of energy (often but not always thermal energy) from a body of
higher temperature to a body of lower temperature.
5. The concepts of kinetic and potential energy can be used to explain physical
processes such as fusion (melting), solidification (freezing), vaporization (boiling,
evaporation), condensation, sublimation, and deposition.
6. Processes that are exothermic give off heat energy. This typically causes the
surrounding environment to become warmer.
7. Processes that are endothermic absorb energy. This typically causes the
surrounding environment to become colder.
42
Topic 6A: Heat & Temperature Outline
4. Heat of fusion (Hf) is the energy needed to convert one gram of a substance from
solid to liquid.
6. Specific heat (C) is the energy required to raise one gram of a substance 1 degree
(Celcius or Kelvin).
The specific heat of liquid water is 1 cal/g*J or 4.2 J/g*K.
43
7. The three phases of matter are solid, liquid and gas. Each has its own properties.
Solids have a constant volume and shape. Particles are held in a rigid, crystalline
structure.Liquids have a constant volume but a changing shape. Particles are mobile but
still held
together by strong attraction.
Gases have no set volume or shape. They will completely fill any closed contained.
Particleshave largely broken free of the forces holding them together.
The phase a substance is in is dependent on the temperature. Melting points and
boilingpoints are on Table S (in Kelvin degrees).
8. Phase changes are a type of physical change. If they are changes that involve
heat being absorbed, they are endothermic changes.
Endothermic phase changes are melting, boiling, evaporating and subliming (sl).
Opposite type of phase changes (freezing, condensing, depositing) are exothermic.
9. A heating curve (or cooling curve) traces the changes in temperature of a
substance as it changes from solid to liquid to gas (or gas to liquid to solid).
When the substance undergoes a phase change, there is no change in temperature. The line
“flattens” until the phase change is complete.
When a phase change is occurring, the potential energy of the substance changes while
kinetic energy remains the same.
As temperature increases, kinetic energy increases.
10. The amount of heat involved in some chemical changes is shown on Table I,
called “heat of reaction” or ∆H.
If the value is negative, the reaction is exothermic.
This can be expressed as a potential energy diagram.
If the energy is written into the equation, and is on the reactants side, the reaction is
endothermic.
∆ H is the difference between the energy stored in the products (PE) and the
potential energy of the reactants.
11. Breaking bonds is ALWAYS endothermic, and forming bonds is ALWAYS
exothermic.
I + II2Bond is forming, I atoms are become stable by bonding, so they
releaseenergy (Exo)
H2 H + H Bond is breaking, requires energy in order to put atoms in unbonded
state (endo)
44
Heat and Temperature – Cut from Jan 2007 – Jan 2008 Exams
1. Given the balanced equation: 3. Which term refers to the difference between
I + I → I2 the potential energy of the products and the
Which statement describes the process represented by potential energy of the reactants for any chemical
this equation? change?
(1) A bond is formed as energy is absorbed. (1) heat of deposition
(2) A bond is formed and energy is released. (2) heat of fusion
(3) A bond is broken as energy is absorbed. (3) heat of reaction
(4) A bond is broken and energy is released. (4) heat of vaporization
7. Determine the total amount of energy released when 2.50 moles of propane is completely reacted
with oxygen. [1]
45
8. Given the balanced equation representing a reaction: N2(g) + O2(g) + 182.6 kJ → 2NO(g)
Draw a potential energy diagram for this reaction. [1]
A 5.00-gram sample of liquid ammonia is originally at 210. K. The diagram of the partial heating
curve below represents the vaporization of the sample of ammonia at standard pressure due to the addition of
heat. The heat is not added at a constant rate.
Some physical constants for ammonia are shown in the data table below.
9. Calculate the total heat absorbed by the 5.00-gram sample of ammonia during time interval AB. Your
response must include both a correct numerical setup and the calculated result. [2]
10. Describe what is happening to both the potential energy and the average kinetic energy of the molecules in
the ammonia sample during time interval BC. Your response must include both potential energy and average
kinetic energy. [1]
11. Determine the total amount of heat required to vaporize this 5.00-gram sample of ammonia at its boiling
point. [1]
46
Base your answers to questions 12 through 14 on the information below.
A 100.0-gram sample of NaCl(s) has an initial temperature of 0°C. A chemist measures the
temperature of the sample as it is heated. Heat is not added at a constant rate. The heating curve for the sample
is shown below.
12. Determine the temperature range over which the entire NaCl sample is a liquid. [1]
13. Identify one line segment on the curve where the average kinetic energy of the particles of the NaCl
sample is changing. [1]
14. Identify one line segment on the curve where the NaCl sample is in a single phase and capable of
conducting electricity. [1]
47
Base your answers to questions 15 and 16 on the information below.
A student performed an experiment to determine the total amount of energy stored in a peanut. The
accepted value for the energy content of a peanut is 30.2 kilojoules per gram. The student measured 100.0
grams of water into a metal can and placed the can on a ring stand, as shown in the diagram below. The peanut
was attached to a wire suspended under the can. The initial temperature of the water was recorded as 22.0°C.
The peanut was ignited and allowed to burn. When the peanut finished burning, the final water temperature
was recorded as 57.0°C. The student’s experimental value for the energy content of this peanut was 25.9
kilojoules per gram.
15. Calculate the total amount of heat absorbed by the water. Your response must include both a correct
numerical setup and the calculated result. [2]
16. Determine the student’s percent error for the energy content of this peanut. [1]
48
Base your answers to questions 17 through 20 on the information below.
The temperature of a sample of a substance is increased from 20.°C to 160.°C as the sample absorbs
heat at a constant rate of 15 kilojoules per minute at standard pressure. The graph below represents the
relationship between temperature and time as the sample is heated.
18. Draw at least nine particles in the box, showing the correct particle arrangement of this sample during the
first minute of heating. [1]
19. What is the total time this sample is in the liquid phase, only? [ 1]
20. Determine the total amount of heat required to completely melt this sample at its melting point. [ 1]
49
Answer Key
p. 5-7 Matter
1. 3 8. 3
2. 2 9. 2
3. 3 10. 3
4. 1 11. 1
5. 4 12. 4
6. 3 13. 3
7. 1 14. 4
15. density of neon gas = 0.827 grams/Liter
16. one physical property is the melting of the wax
17. one indication of chemical change is the statement “the candle burns”… burning (combustion) is
always a chemical change
18. for the solid diagram, show all the particles touching and with a regular pattern to their
arrangement, for the liquid, still show all the particles touching, but do not show them having a
regular arrangement
19. Sample 3
20. Two molecules can be made, leaving one particle of “y” and 4 particles of “z” left over.
21. This symbol does not represent a compound because only one type of element is shown.
83
20. Si-29 contains 15 neutrons
21. Atomic Mass calculation set up:
(27.98 amu) x 0.9222
(28.98 amu) x 0.0469
+ (29.97 amu) x 0.0309
84
p. 26-27 Equations & Stoichiometry
1. 4 8. 3
2. 4 9. 2
3. 4 10. 3
4. 2 11. 1
5. 2 12. 3
6. 2 13. 3
7. 2 14. 4
22. The reaction is endothermic because “heat” is written on the reactants side.
23. To do this question, use the mole ratio between the two substances:
NaHCO3: CO2
2: 1
7: x 2x = 7 x = 3.5 moles of CO2
27. The IUPAC (systematic) name for this compound is Iron III oxide. The roman numeral 3 is
needed because Iron ions can be charged +2 or +3, and is chosen in this case because the formula
includes iron with the +3 charge.
28. moles = 19 / 95 = 0.20 moles
29. The type of ORGANIC reaction is substitution.
30. One advantage of using ozone is that it is safer to use. Another is that it is more environmentally
friendly.
31. Three significant figures are shown.
32. % water mass = ( 0.76 g of water/ 2.13 g of hydrate) x 100% = 35.7%
33. The crucible containing the sample must be heated until a constant mass is achieved in order to
insure that all the water has been driven out of the hydrate
85
p. 35-36 Bonding
1. 3 7. 1
2. 2 8. 1
3. 1 9. 2
4. 1 10. 2
5. 2 11. 4
6. 3 12. 4
p. 45-49 Heat and Temperature Cut from Jan 2007-Jan 2008 Exams
1. 2
2. 1
3. 3
4. 4
5. 2
6. The diagram should look like
7. According to the chemical equation, reacting 1 mole of propane (C3H8) releases 2219.2 kJ , so 2.5
moles x 2219.2 kJ = 5548.0 kJ
86
8. The diagram should look like
q = 707 J
9. average KE = remains the same since Temperature is not changing along BC.
PE = increasing as the sample vaporizes
10. Use the q = mHv equation from Table T
q = 6850 J
11. The entire sample of NaCl is in the liquid phase along segment CD, where the temperature
changes from 801 to 14650C, so the range over which the entire sample is in the liquid phase
is (1465-801) = 664oC
12. The average KE is changing anywhere where the temperature is changing, so correct answers are
segments AB or CD
13. The segment where NaCl is in one phase and able to conduct electricity would be where it is all
and only in the liquid phase (it then has mobile ions capable of conducting electricity). So the
segment CD applies.
14. Use the q = mc T equation from Table T
Obtain the specific heat of water value from Table B, and the change in temperature value
from the paragraph above the picture (22oC to 57oC)
q = (100.0 g) ( 4.18 J/g oC)
(35oC) q = 14630 J