Chemistry Honors
1.
Semester 1 Final Review
Classify each of the following as a physical or chemical change:
ice melting
physical
paper burning
chemical
Acid reacting with oxygen
chemical
gas under pressure
physical
liquid evaporating
physical
Magnesium being heated
chemical
2. What is a chemical change and what is a physical change?
A chemical change involves a rearrangement of the elements in a substance to form substances with
different physical properties. A physical change may change the state of a substance will not change
the composition of that substance.
3. Name three characteristics of most nonmetals.
Brittle, low electrical and thermal conductivity, low boiling point
4. Name three characteristics of metals.
Malleable, ductile, good conductors of heat and electricity, luster.
5. Name three characteristics of most metalloids.
Semiconductors of electricity, solid at room temperature, less malleable that metals.
6. Name two properties of noble gases. Generally unreactive, gas at room temperature.
7. Consider the burning of gasoline and the evaporation of gasoline. Which process represents a chemical
change and which represents a physical change? Give a reason for you answer.
The burning of gasoline represents a chemical change because the gasoline is being changed into
substances with different identities, Evaporation involves a physical change; the identity of gasoline
remains unchanged.
8. Report the number of significant figs in each of the following values.
0.002037g
350. J
64 mL
1.3 x 102 cm
1.30 x 102 cm
34050
9. Round the following top 3 significant figures:
22.77g
14.62 m
9.53052 L
87.55 cm
30.25 g
4
3
2
2
3
4
22.8g
14.6 m
9.53 L
87.6 cm
30.3
10. Three students were asked to determine the volume of a liquid by a method of their choosing. Each did
three trials.. The table below shows results. ( The actual volume is 24.8 mL.)
Trial (mL)
Trial 2 (mL)
Trial 3 (mL)
Student A
24.2
24.6
24.0
Student B
24.2
24.3
24.3
Student C
24.7
24.8
24.7
a.) Which students measurements showed greatest accuracy?
Student C
b.) Which students measurements showed the greatest precision? Student C and B (correction in table)
11. A single atom of platinum has a mass of 3.25 X 10 –22 g. What is the mass of 6.0 X 1023 platinum atoms?
2.0 X 102 g.
12. A sample thought to be pure lead occupies a volume of 15.0 mL and has a mass of 160.0g.
a.) determine its density.
10.7 g/mL
13. 84.01g of baking soda, NaHCO3,, always contains 22.99g of sodium, 1.01g of hydrogen, 12.01g of carbon, and
48.00g of oxygen. What percentage by mass of each of these elements is present in baking soda?
Element
sodium
hydrogen
Percentage
27.37%__
1.20%
Element
carbon
oxygen
Percentage
14.30%__
57.14%__
a. Which law do these data illustrate? The law of definite proportions
14. The smallest unit of an element hat can exist either alone or in combination with atoms of the same or
different element is the __atom__
15. . A positively charged particle found in the nucleus is called a(n) __proton__
16. . A nuclear particle that has no electrical charge is called a(n) __neutron__
17. . The subatomic particles that are least massive and most massive, respectively , are the
__electron__ and __neutron__
18. . A cathode ray produced in a gas-filled tube moves away from a negative field, such as one produced by a
magnet. When a paddle wheel is installed inside the tube, the wheel moves down the tube in the same direction as
the cathode ray. What properties of electrons do these two phenomena illustrate? Who did this experiment?
- __electrons has charge and mass.__ Thompson
19 . How would the electrons produced in a cathode ray tube filled with neon gas compare with the electrons
produced in a cathode ray tube filled with chlorine gas?
- __The electrons produced from neon gas and chlorine gas would behave in the same way because
electrons do not differ from element to element.
20. Explain the difference between the mass number and the atomic number of a nuclide.
- Mass number is the total number of protons and neutrons in the nucleus of an isotope. Atomic number
is the total number of protons only in the nucleus of each atom of an element.
21.
How many particles are in 1 mole of carbon? 1 mol of lithium? 1 mol of eggs? Will 1 mol of each of these
substances have the same mass? There are 6.022 x 1023 particles in 1 mol of each of these substances. A
mole of each substance will not have the same mass.
22.
As the atomic masses of the elements in the periodic table increase, what happens to each of the
following:
a. the number of protons__increases__
b. the number of electrons__increases__
c. the number of atoms in 1 mol of each element stays the same__
23. What is the mass in grams of 2 mol of oxygen atoms? _32.00g__
24 The element boron, B, has an atomic mass of 10.81 amu according to the periodic table. However, no single
atom of boron has a mass of exactly 10.81amu. How can you explain this difference? The periodic table reports
the average atomic mass, which is a weighted average of all isotopes of boron.
25. How did the outcome of Rutherford’s gold foil experiment indicate the existence of a nucleus?
- The particles rebounded, and therefore must have hit a dense bundle of matter. Because such a small
percentage of particles were redirected he reasoned that this clump of matter, called the nucleus,
must occupy only a small fraction of the atoms’ total space.
26. A certain element exists as three natural isotopes as shown in the table below:
Isotope
Mass (amu)
Percent natural abundance
Mass number
1
19.99244
90.51
20
2
20.99295
0.27
21
3
21.99128
9.22
22
Calculate the average atomic mass of this element__20.17945 amu __
27. What is the difference between the ground state and the excited state of an atom (its electrons) ?
- In order for an electron to be ejected from a metal surface, the electron must be struck by a single
photon with at least the minimum energy needed to knock the electron loose.
28. The wavelength of light is 310nm. Calculate its frequency. _9.7 x 1014 Hz
29. What is the wavelength of electromagnetic radiation if its frequency is 3.2x 10-2 Hz.? 9.4 x 109m
30. Compare and contrast Hund’s rule with the Pauli exclusion principle.
31. Explain the conditions under which the following orbital notation for helium is possible:
↑
1s
↑
This orbital notation is possible if an electron in helium’s orbital has been excited.
2s
32. Write the electron configuration and orbital diagrams for each of the following atoms:
a Phosphorus:
1s22s22p63s22p3
b
↑↓ ↑b ↑b ↑↓ ↑↓ ↑↓ ↑ ↑ ↑
1s 2 s 2 px 2 py 2 pz 3s 3 px 3 py 3 pz
Nitrogen 1s22s22p3 ↑↓ ↑↓ ↑ ↑ ↑
1s 2 s 2 px 2 py 2 pz
33.
34.
Write the electron configuration of the following atoms:
a. Carbon _____________ _1s22s22p2
b copper_______________________________ 1s22s22p63p64s13d10
What would be the wavelength of light that has a frequency of 3 x 10-4 Hz in a vacuum? 1 x 1012m
35
When an electron is added to a neutral atom, energy is ____C_
(a) always absorbed
(c)either absorbed or released
(d) always released
(d)burned away
36.
The energy required to remove an electron from an atom is the atom’s ____D
(a) electron affinity
(c)electro negativity
(b)electron energy
(d) ionization energy
37. Moving from left to right across a period on the periodic table.
a. electron affinity values tend to become __(more negative or more positive) _more negative __
b. ionization energy values tend to become __(larger or smaller)
_larger _________
c. atomic radii tend to become __(larger or smaller)
_smaller ________
38.
39
a. Name the halogen with the least – negative electron affinity. _At _
b. Name the alkali metal with the highest ionization energy Li _
c. Name the element in Period 3 with the smallest atomic radius. _Ar _
d. Name the Group 14 element with the largest electro negativity. _C _
a. Compare the size of the radius of a positive ion to its neutral atom.
In positive ions, the radius is smaller than its corresponding neutral atom.
b. Compare the size of the radius of a negative ion to its neutral atom.
In negative ions, the radius is larger than its corresponding neutral atom
40. For metals and nonmetals, which tend to form positive ions? Which tend to form negative ions?
Metals tend to form positive ions; nonmetals tend to form negative ions
41. Explain the role of valence electrons in the formation of chemical compounds.
Valence electrons are the electrons most subject to the influence of nearby atoms or
ions. They are the electrons available to be lost, gained, or shared in the formation
of chemical compounds
42. Consider a neutral atom with 53 protons and 74 neutrons to answer the following questions.
a. What is atomic number? _________ 53
b. What is the mass in amus?__________ 127 amu
c. Is the element’s position in a modern periodic table determined by its atomic number or by its atomic mass?
atomic number
43. Consider an element whose outermost electron configuration is 3d104s24px
a. To which period does the element belong? ______ _Period 4
b. If it is a halogen, what is the value of x?
_5 ____________
c. The group number will equal (10+2+x). True or false? _ true __________
44.. a. Metalloids are found in which block, s,p,d, or f? __p__
b. The hardest, densest metals are found in which block, s, p, d, or f? __d__
45. Referring only to the periodic table , answer the following questions on periodic trends.
a. Which has the larger radius, Al or ln? _ ln __
b. Which has the larger radius, Se or Ca? _Ca __
c. Which has a larger radius, Ca or Ca2++?_Ca __
d. Which has greater ionization energies as a class, metals or nonmetals? _nonmetals __
e. Which has the greater ionization energy, As or Cl? _Cl __
f. an element with a large negative electron affinity is most likely to form a positive ion, a negative ion, or
a neutral atom? _negative ion __
46. Explain why ionic bonds tend to form between metals and nonmetals.
Ionic bonds form between elements with electronegativity differences greater than 2.0. Metals tend to
have low electronegativity while nonmetals have high electronegativity. Metals will more easily give up
electrons to nonmetals, whose atoms are more likely to accept electrons.
47. Explain why the formula of glucose is written as C6H12O6 and not as CH2O.
The formula of glucose is written as C6H12O6 and not as CH2O because a molecule of glucose contains 6
carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms
48. What do the letters VSEPR stand for? Explain what the theory says and how it is used. VSEPR stands for
valence-shell electron pair repulsion. The theory states that, in small molecules, pairs of valence electrons
are arranged as far apart as possible. The theory is used to account for molecular shapes.
49. Explain why silicon tetrafluoride (SiF4) has a square structural formula but a tetrahedral arrangement in
space. The structural formula is a two-dimensional representation and does not take into account the
availability of the space above and below the plane of the molecule. If the atoms were arranged in a planar
square, the bond angles would be 90°, but in the tetrahedral orientation (the correct one) the angles are
109.5°, a wider and more desirable separation of electrons.
50. A scientist is plotting the location of the bonding electrons in vaporized beryllium hydride (BeH2) molecules.
On the basis of beryllium's electron configuration, she expects the graph to show a spherical s orbital that
beryllium shares with one hydrogen and a dumbbell-shaped p orbital that beryllium shares with the other
hydrogen. She also expects one bond to be shorter than the other because of the different kinds of orbitals.
Describe what the scientist will actually find. Hybridization of the s and the p orbitals of beryllium occurs so
that two sp orbitals result. Thus, a dumbbell-shaped probability distribution will be observed rather than
the shapes characteristic of separate s and p orbitals. Thus, the two bonds will be equal in length because
of the equivalence of the two sp orbitals
51. How many valence electrons does the sulfide ion have? The sulfur atom? The sulfide ion has 8 valence
electrons and the sulfur atom has 6 valence electrons
52
How many valence electrons are in the polyatomic ion HPO42–? phosphate ion has 32 valence electrons
(1 from hydrogen, 5 from the phosphorous, 24 from the four oxygen atoms, and 2 from the 2– charge).
53.
How many electrons are shared among carbon atoms in the compound ethanol, CH3CH2OH? Two electrons
are shared among the carbon atoms in ethanol
54 Sulfur is the central atom in sulfur dichloride, SCl2. What is the shape of this molecule? What are the
approximate bond angles?
In SCl2, sulfur is connected to each chlorine by a single bond and also has two
unshared electron pairs. According to VSEPR, the shape should be bent, and the unshared pairs should
reduce the bond angle to about 105°.
55. Describe the shape for a molecule of silane, SiH4. In SiH4, silicon is connected to each hydrogen atom by
a single bond and has no unshared pairs. According to VSEPR, the molecule should be tetrahedral and the
bond angles about 109.5°.
56. What is the shape of a molecule of nitrosyl chloride, NOCl, given that nitrogen is the central atom? What are
the bond angles? The nitrogen atom forms 3 bonds., the oxygen :two bonds, & chlorine atom one bond. N is
central atom, there must be a single bond between nitrogen and chlorine and a double bond between
nitrogen and oxygen. Nitrogen also has an unshared electron pair. Keeping the oxygen, chlorine, and
unshared pair of electrons a maximum distance apart requires a trigonal planar electron-pair geometry and
a bent shape. The bond angle is approximately 120°.
57. A molecule has boron as its central atom and two chlorine atoms and one iodine atom, each connected by a
single bond to boron. Determine if the molecule is polar, given the following electronegativities: B, 2.0; I, 2.5; and
Cl, 3.0. Each bond is polar, with boron at the positive end of each. The molecule is trigonal planar, but in
this molecule the polarities do not cancel because the chlorines are more electronegative than the iodine.
The molecule as a whole is polar, with the chlorine end having a slightly negative charge.
58. Fill in the following table with the correct name or formula:
Name
Beryllium Chloride
Iodine Heptafluoride
Chromium (III) nitrite
Chlorine Trifluoride
Diphosphorus trioxide
Calcium acetate
Binary , Ternary or Molecular
Binary
Molecular
Ternary
Molecular
Molecular
Ternary
Formula
BeCl2
IF7
Cr(NO2)3
ClF3
P2O3
Ca(C2H3O2)2
59. A compound has the empirical formula NO2. Its molecular formula could be: __a___
a. NO2.
b. N2O. c. N4O2. d. N4O4.
60. A compound consists of 85 percent silver and 15 percent fluorine by mass. What is its empirical formula?
AgF