Download Unit #1 Review – Answers Chapter 1 Review – p. 62 #1, 4

Unit #1 Review – Answers
Chapter 1 Review – p. 62 #1, 4 (d,e,g,h), 8, 13, 14, 15, 17, 19, 20, 21.
1. What are the major differences between metals and nonmetals?
Metals are malleable, ductile and conductors of electricity. They are shiny. Nonmetals
are generally nonconductors on electricity in their solid form. At room temperature, they
are mostly gases or solids. Solid nonmetals are brittle.
4. Use the periodic table to answer the following:
d) What are the atomic numbers of hydrogen, oxygen, aluminum, silicon,
chlorine, and copper? The atomic number of hydrogen is 1, of oxygen is 8, or
aluminum is 13, of silicon is 14, of chlorine is 17, and of copper is 29.
e) Identify an element with six electrons in the outer energy level. Oxygen has 6
electrons in its outer energy level.
g) Identify an element that would tend to lose two electrons. Beryllium would tend
to lose 2 electrons.
h) Identify an element that would tend to gain one electron. Fluorine would tend to
gain 1 electron.
8. What is unusual about the atomic structure of hydrogen compared with other
elements? Hydrogen has only one electron in the first energy level. All other elements
have two electrons in the first energy level.
13. From a representative element’s position in the periodic table, how would you
determine each of the following? This example uses the representative element
chlorine:
a) number of protons: The atomic number of chlorine is 17. Therefore, there
are 17 protons.
b) number of electrons: The number of protons equals the number of electrons.
Therefore, chlorine has 17 electrons.
c) number of valence electrons: Chlorine is in group 17, which means it has 7
valence electrons.
d) number of occupied energy levels: Chlorine is in Period 3, so it has three
occupied energy levels.
14. List the number of protons, electrons, and valence electrons in each of the
following atoms: a) magnesium: 12 protons, 12 electrons, 2 valence electrons.
b) aluminum: 13 protons, 13 electrons, 3 valence electrons.
c) iodine: 53 protons, 53 electrons, 7 valence electrons.
15. Write the chemical name and symbol corresponding to each of the following
theoretical descriptions: a) 11 protons and 10 electrons: sodium ion, Na+
b) 18 electrons and a net charge of 3- Phosphorus ion, P3c) 16 protons and 18 electrons: sulfur ion, S217. Determine the number of protons, electrons, and neutrons present in an atom
of each of the following isotopes: a) calcium-42: 20 protons, 20 electrons, 22
neutrons
b) strontium-90: 38 protons, 38 electrons, 52 neutrons
c) cesium-137: 55 protons, 55 electrons, 82 neutrons
d) iron-59: 26 protons, 26 electrons, 33 neutrons
e) sodium-24: 11 protons, 11 electrons, 13 neutrons
19. How does chemical reactivity vary:
a) Amount elements in group 1 and 2? Within group 1 and 2 elements,
reactivity increases moving down the group.
b) Among the elements within groups 16 and 17? within group 16 and 17
elements, reactivity decreases moving down the group.
c) within period 3? Within a period, reactivity tends to be high in Group 1
metals, lower toward the middle of the table, and increase to a maximum in
Group 17 nonmetals.
d) Within group 18. Within group 18, elements are stable and extremely
unreactive.
20. Describe the trends in the periodic table for each of the following atomic
properties, and give a theoretical explanation for each trend:
a) Atomic radii decrease as you move from left to right across each period. From left
to right across a period, the nuclear charge increases while the shielding effect provided
by the non-valence electrons remains the same. As a result of the increasing nuclear
charge, the electrons are more strongly attracted to the nucleus, pulling them closer to
the nucleus and decreasing the size of the atom.
The atomic radii increase as you move down a group. As you move down, there
is an increasing number of energy levels that are filled with electrons. This increases the
shielding effect, decreasing nuclear attraction, and increasing the size of the atom.
b) First ionization energies generally increase as you move from left to right across a
period. From left to right across a period, the nuclear charge increases while the
shielding effect of the non-valence electrons remains the same. As a result the valence
electrons are more strongly attracted to the nucleus, and more energy is required to
remove an electron from the atom.
First ionization energies generally decrease as you move down a group. As you
move down a group, there is an increase in the size of the atoms. As the atomic radius
increases, the distance between the valence electrons and the nucleus also increases.
The non-valence electrons between the nucleus and the valence electrons produce a
shielding effect. As a result, the attraction between the valence electrons and the
nucleus becomes weaker, so less energy is required to remove an electron from the
atom.
c) Electronegativity generally increases as you move from left to right across a period,
as the attraction of the nucleus for any new electrons increases.
Electronegativity decreases as you move down a group, as new electrons are
shielded from the nucleus by increasing numbers of shells of electrons.
d) Electron affinity increases as you move from left to right across a period. Smaller
atoms with larger nuclei have a stronger attraction to new electrons.
Electron affinity decreases as you move down a group. Larger atoms with more
distant and more shielded nuclei have a smaller attraction for new electrons.
Chapter 2 Review – p. 103 #2,3,5,9,10,16(a,b,c,f), 19(a,b,[don’t do c])
2. How does an ionic bond differ from a covalent bond? An ionic bond happens
when one more valence electrons are transferred from a metal atom to a nonmetal
atom. The metal atom becomes a positive ion and the nonmetal atom a negative ion. An
ionic bond is the electrostatic attraction between positive and negative ions in a
compound.
A covalent bond arises from the simultaneous attraction of two nuclei for a
shared pair of electrons. These occur in molecular compounds where the electrons are
shared between two nonmetal atoms and hold the atoms together.
3. a) Briefly summarize and explain the properties of ionic and molecular
compounds. Ionic – solid at room temperature, hard, high melting points and form
solutions that conduct electricity. The properties are due to the strong ionic bonds,
simultaneous forces of attraction between the positive and negative ions which hold the
ions firmly in a rigid structure.
Molecular – At room temperature, may be solids, liquids, or gases, and are soft,
waxy, and flexible. Covalent bonds between the atoms are strong enough (not as strong
as ionic). However, the intermolecular forces in molecular compounds are weaker in
comparison – low melting point and boiling point.
b) Explain why electrical conductivity is generally a suitable test for ionic
compounds. Ionic compounds form solutions that conduct electricity. Because the ionic
bonds often break down in water, the resulting ions are free to move in solution and
conduct electricity. Most molecular compounds from solutions that do not conduct
electricity.
5. For each chemical bond or force listed below, indicate which types of entities
are involved.
a) covalent bond: two nonmetal atoms that are sharing a pair of electrons.
b) dipole-dipole: oppositely charged ends of polar molecules.
c) hydrogen bonds: a positive hydrogen atom of one molecule and a highly
electronegative atom (F, O, or N) in another molecule.
d) ionic bond: a positively charged ion of a metal and a negatively charged ion of a
nonmetal.
9. Draw lewis symbols (electron dot diagrams) to represent a single atom of each
of the following elements: 10. For each of the following substances, predict whether the bonds present are
ionic, covalent, or polar covalent:
a) I2 (s) :covalent, b) SO2 (g): polar covalent, c) OCl2 (g): polar covalent, d)Fe2O3(s): ionic,
e) KBr(s): ionic, f) SrO(s): ionic
16. Write the chemical formula for each of the following substances:
a) magnesium bromide: MgBr2
b) carbon disulfide: CS2
c) mercury(II) nitrite: Hg(NO2)2
f) silver carbonate: Ag2CO3
19. Give the IUPAC names of the following substances:
a) NaCl: sodium chloride,
b) P2O3 (s): diphosphorus trioxide,
c) HNO3(aq) : nitric acid