HILLSBOROUGH
COMMUNITY COLLEGE
WorldWideWolfe
Dr. Drew
H. Wolfe
ATOMS,
ELEMENTS, and the PERIODIC TABLE
Section | 3.2
(Electrons in Atoms) | 3.3 (Periodic Properties)
| Summary | Review
Exercise | Answers | Bottom
| Return to WWWolfe |
Elements are composed of small particles called atoms. Because atoms are tiny and their substructure cannot be directly observed, a theory is required to describe their properties and behavior. Our modern theory of atoms, called quantum theory, has evolved over more than 100 years of research.
Protons, Neutrons, and Electrons
Atoms are composed of protons, neutrons, and electrons. Protons, p+, and neutrons, no, have approximately the same mass (1.67 x 10-24 g). An electron, e-, only has a mass of 9.11 x 10-28 g, about 1/1837 that of a proton (or neutron).
Relative Masses of Subatomic Particles
Masses of subatomic particles are often expressed using a relative unit, termed a unified atomic mass unit, u. One unified atomic mass unit, 1 u, is approximately the mass of a proton (or neutron). The mass of an electron using this scale is 0.000549 u.
Charges of Subatomic Particles
Particles can have a positive (+), negative (-), or no net charge (neutral). Electrons and protons carry the smallest elementary unit of charge found in matter. Electrons are negatively charged (1-). Protons possess have the same magnitude of charge as electrons, but protons possess a positive charge (1+). Neutrons, as the name implies, are electrically neutral particles.
Behavior of Charged Bodies
Particles with the same electric charge repel each other. Objects with unlike charges attract each other. Two electrons or two protons in close proximity repel each other. Unlike charged particles (+ and -), when brought close together, attract each other. The force of electric attraction or repulsion is inversely related to the square of the distance that separates the particles.
The Nucleus
Protons and neutrons are located in a very small region of the atom called the nucleus (plural, nuclei). Most nuclei have diameters roughly 10-6 nm (1 nm = 1 x 10-9 m). Diameters of whole atoms are many times (100,000 x) larger than those of nuclei, ranging from 0.1 to 0.5 nm. Electrons populate the relatively vast space outside the nucleus.
Nuclear Properties of Atoms--Atomic Number
The atomic number, Z, of an atom equals the number of positive charges (number of protons) found in the nucleus. Atoms are arranged in the periodic table in order of increasing atomic number. The atomic number of an element is the integer value found in each box with the atomic mass.
The atomic numbers of H, He, and Li are 1, 2 and 3, respectively. What does the atomic number indicate about each of these atoms?
Solution 3.1
The atomic number is the number of protons in the nucleus of an atom.
Therefore, H, He, and Li have one, two, and three protons, respectively,
in their nuclei.
Nuclear Properties of Atoms--Mass Number
The mass number, A, of an atom equals the number of protons plus neutrons in the nucleus.
Mass numbers are just numbers and are not masses.
Atomic Symbols
To express the atomic number and mass number of an atom, atomic symbols are written as follows
in which X is the symbol of the atom, A is the mass number, and Z is the atomic number.
Write the symbol for the atom that has 36 p+ and 48 no.
Solution 3.2
The periodic table shows that the element with 36 protons (Z = 36) is Kr. To obtain the mass number, A, add the number of protons, 36 p+, to the number of neutrons, 48 no.
A = 36 p+ + 48 no = 84
Hence, the symbol for this atom is 8436Kr.
The Number of Neutrons in an Atom
To obtain the number of neutrons in a nucleus, given the mass number and atomic number, subtract the atomic number (number of p+) from the mass number (number of p+ plus no).
Number of no = mass number - atomic number = (number of p+ + no) - number of p+
Mass numbers are not found on the periodic table.
What is the nuclear composition of 6429Cu?
Solution 3.3
The symbol shows that the atomic number, Z, is 29. Therefore, 29 protons are in this atom. To obtain the number of neutrons, subtract the atomic number, 29, from the mass number, 64.
Number of no = A - Z = 64 - 29 = 35 no
6429Cu atoms have 29 protons and 35 neutrons in
their nuclei.
What is the composition of the nucleus of a 22688Ra atom?
Solution 3.4
The mass number of 22688Ra is 226 and its atomic number is 88. Therefore, 88 protons are in the nucleus. The number of neutrons is calculated by subtracting the atomic number from the mass number.
Number of no = A - Z = 226 - 88 = 138 no
22688Ra contains 88 protons and 138 neutrons in its nucleus.
Isotopes
Isotopes are atoms with the same atomic but have different mass numbers. This means that isotopes have the same number of protons but different numbers of neutrons. For example, three isotopes of hydrogen are known to exist. 1H (protium) and 2H (deuterium) are found in natural samples of hydrogen, and 3H (tritium) is a radioactive isotope that is not found in nature. Each of these isotopes has one proton in the nucleus, but differ in the number of neutrons (0, 1, and 2 no).
23492U, 23592U, and 23892U are three isotopes in naturally occurring samples of uranium. What is the composition of each of the nuclei of these isotopes?
Solution 3.5
Each U nucleus contains 92 protons. These isotopes differ in the number
of neutrons in the nucleus. 234U nuclei contain 142 neutrons,
235U nuclei contain 143 neutrons, and 238U nuclei
contain 146 neutrons.
Masses of Atoms
An atom has a small mass. For example, a H atom has a mass of only 1.67 x 10-24 g. To avoid the inconvenience of working with such small numbers, a relative scale for the masses of individual atoms are used. In this scale, masses of all atoms are expressed relative to the mass of one 126C atom.
Atomic Mass Units
By definition the mass of a 126C atom equals exactly 12 unified atomic mass units, u. Therefore, one unified atomic mass unit, 1 u, is 1/12 the mass of the 126C atom. Because, 126C is composed of six protons and six neutrons, 1 u is about the average mass of a proton and a neutron.
The masses of other atoms are determined relative to the standard, 12C. For example, if an atom is found to have a mass three times that of 12C, its relative mass is 36 u, 3 times 12 u. An atom with a mass one-fourth the mass of 12C is assigned a mass of 3 u, one-fourth times 12 u.
Atomic Mass (Atomic Weight)
The atomic mass of an element is the average mass of its naturally occurring isotopes relative to 12C. Atomic masses, often called atomic weights, are the numbers listed in each box in the periodic table with the chemical symbol and atomic number of an element. Atomic masses of the elements are decimal numbers (e.g., Al, 26.9815; S, 32.06; V, 50.942; etc.) because a large percent of the naturally occurring elements exists as a mixture of isotopes. Each of these isotopes has a different mass. Because a sample of matter has a large quantity of atoms that have different masses, it is convenient to use the average mass of an element's isotopes. This is called the atomic mass.
Atomic Mass of Carbon
Consider the atomic mass of carbon, 12.011 u. Carbon is composed principally of two isotopes: 12C and 13C. In nature, 9889 out of 10,000 carbon atoms are 12C, and 111 are the heavier 13C isotopes. Averaging 9889 particles with a mass of 12.00 u and 111 particles with a mass of 13.00 u gives an average mass of 12.01 u.
In natural samples of the element chlorine, 75.53% is 35Cl with a mass of 34.969 u. The remaining 24.47% is 37Cl, which has a mass of 36.966 u. Calculate the atomic mass of chlorine.
Solution 3.6
To obtain the atomic mass calculate the weighted average of the masses of the two isotopes. To accomplish this, assume that 100 Cl atoms are present. Of these 100 total Cl atoms, 75.53 have a mass of 34.969 u, and 24.47 have a mass of 36.966 u. Therefore, multiply 75.53 times 34.969 u, add that to the product of 24.47 times 36.966 u, and divide by the number of atoms, 100.
Atomic mass (Cl) = [(34.969 x 75.53) + (36.966 x 24.47)]/100 = 35.46
Calculate the atomic mass of boron, using the following data:
| Isotope | Isotope Relative mass, u | Percent abundance, % |
| 10B | 10.013 | 19.70 |
| 11B | 11.009 | 80.30 |
Solution 3.7
The atomic mass of B is obtained by calculating the average mass of its isotopes.
Atomic mass (B) = [10.013 x 19.70) + (11.009 x 80.30)]/100 = 10.81
Electrons
An electron has a small mass with respect to neutrons and protons, a negative charge, and is located outside the nucleus. Because all atoms are electrically neutral, the number of electrons in an atom equals the number of protons (atomic number).
Number of e- = number of p+ = atomic number
It is only necessary to look at the periodic table to find the number of electrons in an atom. For example, H atoms have 1 e-, C atoms have 6 e-, and Ne atoms have 10 e-.
Electron Orbitals
Electrons are elusive particles, ones that cannot be directly observed. Because an electron cannot be seen, the path and exact location at any particular time are unknown. Therefore, scientists can only identify the regions of space where electrons are most probably found. These regions of space are called orbitals. An orbital is a region of space around the nucleus where a high probability exists of finding electrons. The maximum number of electrons that an orbital can hold is two electrons. Orbitals can be empty (0 e-), half-filled (1 e-), or filled (2 e-).
Electron Energy Levels and Shells
A shell is a collection of orbitals at approximately the same average distance from the nucleus. Electrons closer to the nucleus are in a lower-energy shell and electrons farther from the nucleus are in a higher-energy shell. An electron in the first shell (n = 1) on an average, is closer to the nucleus than an electron that occupies the second shell (n = 2). The average distance from the nucleus of electron in the third shell (n = 3) is farther than one in the second shell (n = 2).
Each shell has a maximum number of orbitals and electrons. Lower-energy shells are closer to the nucleus where less volume is available for the electrons to occupy. Mutual repulsion of electrons limits the number of electrons in a given region. Each higher-energy shell has a greater volume for electrons to populate. The shell, represented by the integer n, can have a maximum of n2 orbitals and 2n2 electrons. The first shell has a n value equal to one, the second has a n value equal to 2, and so on.
What is the maximum number of orbitals and electrons in the first three shells, n = 1 to 3?
Solution 3.8
Te first shell has 12, 1 orbital, and 2 x 12,
2 electrons. The second shell has a maximum of 22, 4 orbitals,
and 2 x 22, 8 electrons. The second shell has a maximum of 32,
9 orbitals, and 2 x 32, 18 electrons.
Subshells
Each shell has one or more subshells. A subshell is composed of orbitals in the same shell that have the same characteristics. The number of subshells within a shell corresponds to the value of n of the shell. Thus, the first shell (n = 1) contains only one subshell--they are the same. The second shell (n = 2) is composed of two subshells, the third shell (n = 3) has three subshells, and so on. Each subshell is denoted by a letter: s, the lowest energy subshell; p, next highest energy subshell; d, higher still; and f, the highest energy subshell of the four.
Orbitals and Subshells
Each subshell can hold a maximum population of orbitals and electrons. All s subshells have a maximum of two electrons because the s subshell is only composed of a single orbital. The p subshell has three orbitals. Thus, six electrons are the maximum number of electrons that can populate the p subshell. Five and seven orbitals are associated with the d and f subshells, which hold a maximum of 10 and 14 electrons, respectively.
Shapes of Orbitals
Subshells are distinguished by the shapes of the orbitals of which they are composed. For example, the shape of the s subshell is spherical. A more complex distribution is found for the p subshell. One of the p orbitals is aligned along the x axis, another is along the y axis, and the third is located along the z axis. Note that the "dumbbell" shape of a p orbital is much different from that of the s orbital. The highest probability of finding electrons in p orbitals is along the axes. Both d and f orbitals have more complex distributions and will not be considered in this discussion.
Aufbau Principle
Electrons fill orbitals starting from the lowest energy orbital (n = 1), and proceed, one electron at a time, filling each lower-energy orbital before filling a higher-energy orbital. The filling of lower-energy orbitals before higher energy orbitals is the aufbau principle.
Electron Arrangements--H Through Be
In all atoms, the lowest energy orbital, the one closest to the nucleus, is the 1s orbital. The first number, 1, refers to the shell where the electron is located, and the s identifies the subshell.
Hydrogen is the simplest atom, with only one electron. Its electron configuration--representation of occupied orbitals--is as follows
The one electron in H is found in the s orbital of the lowest shell.
Helium atoms contain two electrons (Z = 2). Because space is available for another electron in the 1s orbital, both electrons occupy this orbital. The electron configuration of helium is
A 2 is written as a superscript above the s to show that two electrons are in the 1s orbital. The orbital is now filled. Because two electrons is maximum number that can occupy the first shell, the first shell is also filled.
Lithium (Z = 3) is the first atom to have an electron in the second shell. Three electrons are found in Li atoms; the first two occupy the lower shell, 1s, and the remaining one is in the next higher energy 2s orbital. The electron configuration of lithium is
Though the second shell contains two subshells, s and p, the s subshell has a lower energy than the p. Accordingly, the 2s subshell fills before an electron enters the 2p subshell.
Write the electron configuration for Be (Z = 4).
Solution 3.9
The two lowest energy electrons in Be occupy the 1s orbital--the He configuration. The two outer electrons are in the 2s subshell, filling it. The configuration of Be is
Electron Arrangements--B Through Ne
In boron (Z = 5) the first four electrons occupy the same orbitals as the four electrons in Be, 1s2 and 2s2. The fifth electron enters the higher energy 2p subshell. Boron is the first element to possess an electron in the 2p subshell. The electron configuration of boron is
Because the p subshell contains three orbitals and can hold six electrons, the next five atoms on the periodic table, C through Ne, fill the 2p subshell.
A Ne atom is the first to have a completely filled second shell.
The s and p subshells in the third shell fill like those in the second. Write the electron configurations for Na (Z = 11) to Ar (Z = 18).
Solution 3.10
The Periodic Table and Electron Configurations
At this point, it is important to recognize that the periodic table is organized according to the arrangement of electrons in atoms. In the outermost shells of the elements in groups 1 (IA) and 2 (IIA), the s subshell fills. The outer p subshell fills in the elements in groups 13 (IIIA) to 18 (VIIIA). The outer d subshell fills in the transition elements (groups 3 to 12), and the outer f subshell fills in the lanthanide (elements 58 to 71) and actinide series (elements 90 to 103). To write electron configurations, follow the periodic table in order of increasing atomic number. Numbers that denote periods correspond to electron shells, and the Roman numeral at the top of each vertical column give the outer-level (valence) configuration.
The 4s Fills Before the 3d
Following the periodic table readily shows that the 4s fills before the 3d. Potassium, K, is a member of group 1 (IA), all of which have one valence electron in the s subshell. Hence, the electron configurations for K and Ca are as follows.
Instead of writing the entire inner core of electrons, the Ar configuration, often [Ar] is written meaning 1s2 2s2 2p6 3s2 3p6.
Note that the outer-shell electron configuration of K is 4s1. Note that the other members of group 1 (IA) of the periodic table also have an outer s1 electron.
All members of the first group have an outer-shell, commonly called the valence shell, configuration of ns1.
What is the valence electron configuration for the group 2 (IIA) elements? Write the shorthand configuration for these elements.
Solution 3.11
All atoms in group 2 (IIA) have two outer electrons in the s subshell, ns2.
In all cases, atoms listed within a group on the periodic table have the same number of valence electrons.
Electron Configurations--Sc to Kr
Scandium, Sc, is the first atom to have an electron in the 3d orbital.
From Sc to Zn, the 3d subshell fills somewhat irregularly. Zinc is the first atom to have a complete 3d subshell. Its electron configuration is
After the 3d is full, the 4p fills. Gallium, Ga, is the first element with a 4p electron.
The first element to have a filled 4p is Kr. The configuration of Kr is as follows.
Use the periodic table to write the complete electron configuration for tin, Sn (Z = 50).
Solution 3.12
Tin has 50 electrons, is in the fifth period, and is a member of group 14 (IVA). All elements in group 14 (IVA) have a valence configuration of two electrons in the s orbital, s2, and two in the p orbital, p2. To write the electron configuration, follow the periodic table in order of increasing atomic number.
Always check to see that the number of electrons in the electron configuration
equals the atomic number. For this example, the sum of the superscripts
equals 50, the atomic number of Sn.
Lewis Symbols
The outer electron configurations of atoms for the most part determine the properties of elements, especially their chemical properties. Lewis symbols (also called dot formulas) are used to conveniently express the valence configurations of atoms. To write a Lewis symbol:
1. Write the symbol of the atom.
2. Place one dot around the symbol for each electron in the valence shell of the atom.
Always remember that inner-level electrons are never included in Lewis symbols.
Write the Lewis symbols for H and He.
Solution 3.13
Because a hydrogen atom only has one electron (1s1), its Lewis symbol has only one dot next to its symbol.
Helium has two electrons (1s2), so its Lewis symbol shows two dots next to its symbol.
Write the Lewis symbols for F, Cl, Br, and I--members of group 17 (VIIA) of the periodic table.
Solution 3.14
All atoms in a chemical group have the same number of dots around their
symbols as a result of having the same number of valence electrons. Atoms
in group 17 (VIIA) all have seven valence electrons. Thus, the general
Lewis symbol for these atoms has seven dots around the symbol.
Draw the Lewis symbols for each of the following. (a) Al, (b) Rb, (c) Xe
Solution 3.15
(a) Aluminum belongs to group 13 (IIIA) and has an outer electron configuration of 3s2 3p1. Each member of group 13 (IIIA) has three valence electrons (s2p1); therefore, the Lewis symbol of Al has three dots around the symbol for Al.
(b) Rubidium belongs to group 1 (IA) and has a valence configuration of 5s1, one outer-shell electron. Consequently, the Lewis symbol for Rb is Rb.
(c) Xenon belongs to group 18 (VIIIA) and has an outer electron configuration
of 5s2 5p6; thus, eight dots are placed
around its symbol.
3.3 PERIODIC PROPERTIES OF ELEMENTS AND ATOMS
Names of the Chemical Groups
Elements within a group of the periodic table are given chemical group names. Elements in group 1 (IA), except H, are called the alkali metals. Those located in group 2 (IIA) are called the alkaline earth metals. The next 10 columns, groups 3 to 12 (B group elements), are the transition metals. Groups 13 (IIIA) and 14 (IVA) have no unique names but are often called the boron-aluminum and carbon-silicon groups, respectively. Group 16 (VIA) is called the chalcogens. Group 17 (VIIA) contains the halogens, and group 18 (VIIIA) contains the noble gases.
Periodic Law
The periodic law states that the properties of the elements and their atoms are a periodic function of their atomic number. In other words, when the elements are listed in order of their atomic numbers, a regular pattern of chemical properties is observed.
Metal, Nonmetals, and Metalloids
The metals are on the left side of the periodic table. They are found to the left of the bold zigzag line, often called the "staircase," that is found on most periodic tables. On the right side of the "staircase" are the nonmetals. Metals are usually solid silver-gray, high density elements that have rather high melting and boiling points, are good conductors of heat and electricity, and are malleable and ductile. Nonmetals most often possess properties opposite to those of metals. Many nonmetals are liquids and gases, which generally have lower melting points, boiling points, and densities than metals. Nonmetals are insulators--they do not conduct heat or electricity.
The elements that exactly border the "staircase" are the metalloids
(sometimes called semimetals). These elements have intermediate properties,
with properties of both metals and nonmetals.
Classify the period 2 elements as metals, nonmetals, and metalloids.
Solution 3.16
The elements in period 2 to the left of the "staircase," the metals,
are Li and Be. The elements to the right of the "staircase," the nonmetals,
are C, N, O, F, and Ne. The metalloid of the period 2 elements is B. Some
would argue that Be is also a metalloid because it borders the "staircase."
Classify the group 14 (IVA) elements as metals, nonmetals, and metalloids.
Solution 3.17
Carbon is a nonmetal. Silicon and Ge are metalloids, and Sn and Pb are
metals.
Consider the elements Cd, Sn, I, and Xe to answer the following questions. (a) Which of these elements have the highest melting points? (b) Which element does not conduct an electric current? (c) Which element has the lowest density?
Solution 3.18
Of these elements, Cd and Sn are metals and I and Xe are nonmetals.
(a) Metals usually have higher melting points than nonmetals; thus, Cd and Sn have the highest melting points.
(b) Nonmetals are electric insulators; hence, I and Xe are nonconductors.
(c) Of the two nonmetals, I and Xe, I is a solid and Xe is a noble gas;
therefore, Xe has a lower density.
Ion Formation--Cations and Anions
Metals have few valence electrons and tend to lose them during chemical changes. Nonmetals, excluding the noble gases, tend to gain electrons in chemical changes. When atoms lose or gain electrons they form ions. An ion is a charged atom or a charged group of atoms. Positively charged ions, called cations, result when metals lose electrons. Negatively charged atoms, called anions, result when nonmetals gain electrons.
Metals most often release their valence electrons and produce cations that have the same electron configurations as the noble gas that precedes it on the periodic table. These cations are said to be isoelectronic to the noble gas. Nonmetals most often take in enough electrons to produce anions that have a noble gas configuration.
(a) Show what happens when a Na atom loses one electron and forms a Na+ cation. (b) Show what happens when a Ca atom loses two electrons and produces a Ca2+ cation.
Solution 3.19
(a) Sodium has the electron configuration of 1s2 2s2 2p6 3s1. When it loses the 3s1 electron, the resulting configuration for the Na+ cation is 1s2 2s2 2p6, the same configuration as Ne. In other words, it is isoelectronic to Ne.
(b) Calcium has the electron configuration of 1s2 2s2 2p6 3s2 3p6 4s2. When it loses its two 4s2 electrons, the resulting configuration for the Ca2+ cation is 1s2 2s2 2p6 3s2 3p6, the same configuration as Ar.
(a) what happens when a F atom gains one electron and forms a F-, fluoride, anion? (b) What happens when a S atom gains two electrons and forms a S2-, sulfide, anion?
Solution 3.20
(a) Fluorine has the electron configuration of 1s2 2s2 2p5. When it gains one electron, the resulting configuration of the F- anion is 1s2 2s2 2p6, the same configuration as Ne. In other words, it is isoelectronic to Ne.
(b) Sulfur has the electron configuration of 1s2 2s2
2p6 3s2 3p4. When
it gains two electrons, the resulting configuration for the S2-
anion is 1s2 2s2 2p6
3s2 3p6, the same configuration as
Ar.
Ionization Energy
Individual atoms also exhibit periodic trends. An important property of atoms that helps explain ion formation is called ionization energy. Ionization energy, IE, is the minimum amount of energy required to remove the most loosely held electron from a neutral gaseous atom, A(g).
The ionization energy of an atom is a measure of the degree to which the nucleus attracts the most loosely held valence electron, the one most distant from the nucleus. A low ionization energy indicates a smaller attractive force and a high ionization energy indicates a larger attractive force between the nucleus and the most loosely held electron.
Trends in Ionization Energy
Two distinct trends are evident when considering the ionization energies of elements. First, going from left to right across a period (increasing atomic number within a period), the ionization energy generally increases. More energy is required, on the average, to remove the most loosely held valence electron from nonmetals than metals. Alkali metals have the lowest ionization energies, and noble gases have the highest ionization energies within a period.
Second, within a group (going from top to bottom) like the noble gases (group 18), ionization energies decrease with increasing atomic number. Helium has the highest ionization energy (2.37 MJ/mol) within the noble gases, and Rn has the lowest (1.04 MJ/mol). Similar trends are found in all chemical groups.
Overall, elements in the lower left corner of the periodic table have the lowest ionization energies; francium, Fr, is the element with the lowest ionization energy. As we move up and to the right on the periodic table, the ionization energies generally increase; He has the highest ionization energy of all atoms.
Select the element from the following groups that has the highest ionization energy. (a) F, Ne, Cl, Ar, (b) O, S, Se, Te, (c) Na, K, Mg, Ca.
Solution 3.21
(a) Neon has the highest ionization energy because noble gases have the highest ionization energies in a period and the lower-mass noble gases have higher ionization energies than higher-mass noble gases.
(b) Oxygen has the highest ionization of these group 16 elements. Smaller atoms have higher ionization energies than larger ones.
(c) Magnesium has the highest ionization energy of this group. It has
a higher ionization energy than Na within its period and a higher ionization
energy than Ca within its group.
The atom is a small particle composed of a dense nucleus that contains protons and neutrons. Electrons sparsely populate the outer regions of the atom. Protons and neutrons have approximately the same mass. Electrons have an extremely tiny mass. However, electrons possess a full negative charge, equal in magnitude but opposite in sign to the proton. Neutrons have no electric charge.
The atomic number of an atom equals the number of protons in the nucleus of an atom. The mass number equals the number of protons and neutrons in the nucleus. Isotopes are atoms with the same atomic number but have different mass numbers. Atomic mass (atomic weight) is the weighted average mass of the isotopes of an element compared to the 12C isotope.
Electrons are located in regions of space called orbitals. Orbitals are regions around the nucleus where a high probability exists of finding electrons. Orbitals hold a maximum of two electrons. A set of orbitals with the similar characteristics and nearly the same energy is called a subshell. Electrons in their lowest energy states are found in four different subshells; s, p, d, and f. Subshells with similar energies are grouped into shells.
Electrons fill orbitals in atoms, starting with the lowest energy orbital and proceed to higher energy orbitals. Each different atom has its own specific electron configuration. Atoms in the same group on the periodic table have the same subshell configuration. The number at the top of each group corresponds to the number of electrons in the outer shell.
The periodic law states that the properties of the elements are periodic functions of their atomic number. When atoms are arranged in order of increasing atomic number, recurring chemical and physical properties are found. Elements with the same outer electron configuration are in the same group (vertical column). Group members share many common properties, and within most groups regular trends in properties are found. Periods (horizontal rows) contain elements that have outer electrons in the same energy level.
Elements are classified as metals, nonmetals, or metalloids (semimetals), depending on their properties. Metals occupy the left side of the periodic table, and are usually solids with high densities, and high melting and boiling points. They are also good conductors of heat and electricity and tend to lose electrons in chemical changes. Nonmetals, on the right side of the periodic table, have lower average densities, melting points, boiling points, and conductivities than metals. In chemical reactions, nonmetals tend to gain electrons and form anions. Metalloids resemble metals but have some nonmetallic properties.
Ionization energy is the amount of energy required to remove the most loosely held electron from a neutral gaseous atom. Across a period from left to right the ionization energy increases. Within a chemical group, the ionization energy decreases with increasing atomic number.
1. Describe what happens when the following charged particles are brought close to each other. (a) two protons, (b) proton and electron, (c) neutron and electron
2. How many protons, neutrons, and electrons are found in each of the following atoms? (a) 10846Pd, (b) 7032Ge, (c) 4822Ti
3. Write the symbols for the atoms that have the following number of protons and neutrons. (a) p+ = 65, no = 94, (b) p+ = 53, no = 74, (c) p+ = 44, no = 58
4. Europium, Eu, contains two isotopes 151Eu (mass = 150.92 u), and 153Eu (mass = 152.92 u). If 151Eu and 153Eu natural abundances are 47.82% and 52.18%, respectively, calculate the atomic mass of Eu.
5. Use the following mass and natural abundance data to calculate the
atomic mass of thallium, Tl:
| Isotope | Isotope Mass, u | Percent abundance, % |
| 203Tl | 202.97 | 29.50 |
| 205Tl | 204.97 | 70.50 |
6. Write the maximum number of electrons that can be located in the following. (a) an orbital, (b) the d subshell, (c) a Be atom, (d) third shell, (e) the f subshell, (f) a Ca atom, (g) Ca2+
7. For each of the following, identify the region closest to the nucleus. (a) fifth, sixth, or seventh shell, (b) 4s, 4p, or 4d subshell, (c) 3s, 4s, or 5s subshell.
8. Write the complete electron configuration for each of the following atoms. (a) Li, (b) C, (c) Na, (d) As, (e) Rb
9. Write the shorthand notation for the electron configuration for each of the following atoms. (a) Ba, (b) Fr, (c) Br, (d) Pb
10. Write the valence electron configuration for each of the following atoms. (a) K, (b) Sr, (c) Ga, (d) Se, (e) Cl
11. Write the symbols for the elements with the following electron configurations.
(a) 1s2 2s2 2p3
(b) 1s2 2s2 2p6 3s2
(c) 1s2 2s2 2p6 3s2 3p3
(d) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
(e) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6
12. Draw the Lewis symbols for each of the following atoms. (a) B, (b) O, (c) Mg, (d) Pb
13. Identify the atom or atoms with the following characteristics. (a) first group 1 (IA) atom to have a 3s valence electron, (b) group 17 (VIIA) atom first to have a 4p electron, (c) atom with atomic number less than 40 that has electrons in the 4d subshell, (d) group 2 (IIA) atom(s) with occupied f orbitals, (e) atom(s) that only have s electrons.
14. What noble gas configuration would each of the following obtain if they lost or gained the stated number of electrons? (a) Mg, loses 2 e-, (b) P, gains 3 e-, (c) Cl, gains 1 e-, (d) O, gains two e-
15. Write the name of the group on the periodic table to which each of the following elements belongs. (a) Na, (b) Al, (c) O, (d) As, (e) Cu, (f) I, (g) Be
16. What is the valence electron configuration for each of the following groups of elements? (a) alkali metals, (b) halogens, (c) chalcogens, (d) noble gases, (e) alkaline earth metals, (f) nitrogen-phosphorus group elements
17. Classify each of the following as a metal, nonmetal or metalloid. (a) Ar, (b) N, (c) H, (d) Tl, (e) Li, (f) Sr
18. Consider the following properties of hypothetical elements A and B, and classify each as a metal, nonmetal, or metalloid. (a) Element A boils at -195.8oC, has a density of 1.3 g/dm3, and is a colorless gas at room temperature. (b) Element B boils at 3200oC, has a density of 10 g/cm3, and is a good conducting solid.
19. Select the most metallic element in each of the following groups. (a) Be, B, C, N, (b) C, Si, Ge, Sn; (c) As, Sn, Sb, Ge, (d) Zn, S, Si, P
20. How many electrons must be removed from each of the following atoms to yield a noble gas configuration? (a) Cs, (b) Be, (c) Al, (d) Sc, (e) Mg
21. What charge would the most stable ions of the following atoms possess? (a) Sr, (b) P, (c) S, (d) O, (e) N, (f) Rb
22. What noble gases are isoelectronic to the following ions? (a) Ca2+, (b) Al3+, (c) N3-, (d) I-, (e) Rb+
23. From the following groups of atoms predict the atom with the highest
ionization energy. (a) C, Si, Ge, Sn, (b) As, Se, Br, Kr, (c) K, Ca, Rb,
Sr, (d) F, Ne, Cl, Ar
1. (a) repel, (b) attract, (c) no attraction or repulsion
2. (a) p+ = 46, e- = 46, no = 62, (b) p+ = 32, e- = 32, no = 38, (c) p+ = 22, e- = 22, no = 26
3. (a) 15965Tb, (b) 12753I, (c) 10244Ru
4. 152.0
5. 204.4
6. (a) 2 e-, (b) 10 e-, (c) 4 e-, (d) 18 e-, (e) 14e-, (f) 20 e-, (g) 18e-
7. (a) 5th shell, (b) 4s, (c) 3s
8. (a) 1s2 2s1
(b) 1s2 2s2 2p2
(c) 1s2 2s2 2p6 3s1
(d) 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p3
(e) 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s1
9. (a) [Xe] 6s2, (b) [Rn] 7s1, (c) [Ar] 4s2 3d10 4p5, (d) [Xe] 6s2 4f14 5d10 6p2
10. (a) 4s1, (b) 5s2, (c) 4s2 4p1, (d) 4s2 4p4, (e) 3s2 3p5
11. (a) N, (b) Mg, (c) P, (d) Ge, (e) Xe
12. (a) B (three dots), (b) O (six dots), (c) Mg:, (d) :Pb:
13. (a) Na, (b) Br, (c) Y, (d) Ra, (e) H and He
14. (a) Ne, (b) Ar, (c) Ar, (d) Ne
15. (a) alkali metal, (b) boron-aluminum group, (c) chalcogen, (d) nitrogen-phosphorus, (e) transition metal, (f) halogen, (g) alkaline earth metal
16. (a) ns1, (b) ns2 np5, (c) ns2 np4, (d) ns2 np6, (e) ns2, (f) ns2 np3
17. (a) nonmetal, (b) nonmetal, (c) nonmetal, (d) metal, (e) metal, (f) metal
18. (a) nonmetal, (b) metal
19. (a) Be, (b) Sn, (c) Sn, (d) Zn
20. (a) 1 e-, (b) 2 e-, (c) 3 e-, (d) 3 e-, (e) 2 e-
21. (a) 2+, (b) 3-, (c) 2-, (d) 2-, (e) 3-, (f) 2+
22. (a) Ar, (b) Ne, (c) Ne, (d) Xe, (e) Kr
23. (a) C, (b) Kr, (c) Ca, (d) Ne
Section | 3.1
(Atoms) | 3.2 (Electrons in Atoms) | 3.3
(Periodic Properties) | Summary | Review
Exercise | Answers
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