• Patterns within families are not the only patterns on the periodic table. Patterns also exist across periods. Most of the trends in properties across a period in the periodic table are the result of the arrangement of the electrons in the atoms in the periods. Moving from left to right across a period, electrons are being added to the same energy level. At the same time, protons are being added to the nucleus. The addition of protons increases the positive charge of the nucleus. A nucleus with a greater positive charge attracts the electrons in the atom more strongly than does a nucleus with less positive charge and the same number of electron energy levels. This attraction is the basis for many of the observed trends across a period.
• A property that decreases from left to right across a period is atomic size. The alkali metal in a period is the largest atom in the period. For example, sodium (Na) is the largest atom in period 3. The next atom in the period, magnesium (Mg), has one additional proton and one additional electron that is in the same energy level as the outermost electron in sodium. This additional positive charge pulls the electrons in the atom closer to the nucleus. Thus, a magnesium atom is smaller than a sodium atom—a trend that continues across the period.
• Ions form when atoms gain or lose electrons. Ionization energy is the amount of energy it takes to form a positive ion from a neutral atom by removing an electron from it. Alkali metals have relatively low ionization energies because the outer electron is far from the nucleus and is easily removed. Moving from left to right across a period, the outer electrons are attracted more strongly by the nucleus and thus are more difficult to remove. Although there are exceptions in each period, in general, ionization energy increases from left to right across a period.
• Some periodic trends are based on the way that elements react with one another. Atoms tend to gain or lose electrons to achieve a stable electron configuration. While there are other stable configurations, the most stable electron configuration is a complete outer electron energy level. For most atoms in the upper part of the table, this is eight electrons in the outer level. For example, a sodium atom has 11
electrons—two in the first energy level, eight in the second, and one in the third. Sodium has a tendency to lose its single electron, giving it a complete second energy level. Near the other end of the same period, chlorine has seven electrons in its outer energy level. It will most easily gain one electron, thus having a complete third energy level.
• Because of the way atoms react, the reactivity of metals generally decreases from left to right within a period. Conversely, the reactivity of non-metals increases from left to right within a period. Note, this generality does not include the noble gases. These elements already have eight electrons in their outer energy levels. They are already stable and do not have a tendency to react with other atoms.
POSSIBLE MISCONCEPTIONS Identify
• Students may think that all atoms of the same element have the same mass.
Clarify
• All neutral atoms of the same element have the same number of protons and the same number of electrons. However, atoms of the same element can have different numbers of neutrons. For example, approximately three quarters of all chlorine atoms contain 18 neutrons. Most of the rest of chlorine atoms contain 20 neutrons. The chlorine atoms with 18 neutrons have a different mass than those with 20 neutrons. (These
different forms of the same element are called isotopes.)
Ask What They Think Now
• At the end of the lesson, ask, How can atoms of the same element have different masses? Students should realize that when atoms of the same element have different
numbers of neutrons, they have different masses.
TEACHING NOTES Engage
• Tell the class that for centuries, alchemists were people who tried various
methods to make gold out of “baser” metals, such as copper or tin. In fact, modern chemistry in many ways grew out of the efforts of these alchemists. Review the example on page 234 of the Student Book in which copper and tin are combined (in theory) to have their two atomic masses (29 and 50) sum to a total of 79 to match gold. Ask, If you wanted to make gold what other combinations could you use and how would you put them together? Allow the students to discuss the
issue. (Sample answer: nickel (28) and antimony (51))
Explore and Explain
• To understand the use of atomic mass units (u) for atoms rather than the familiar
grams and kilograms, ask, Which units are generally used to measure mass? (sample
answer: gram, kilogram) Why would it be difficult to measure atomic mass in grams or kilograms? (sample answer: The numbers would be too small to work with easily.) Point out that one atom of iron has a mass of 9.29 × 10–23 g. Rather than use such small and awkward numbers, scientists invented the atomic mass unit.
• To understand how atomic masses are computed, present the concept of a weighted average. Ask, If 9 students in a class score 100 % on a test, and the other 1 student scores 90 %, what is the class average? (99 %) Point out that that the atomic mass of
an element is computed using a weighted average in much the same way.
• Refer students to Sample Problem 1: Write a word equation on the board: atomic number + neutrons = atomic mass. Now have volunteers provide values for each part
of the equation. Atomic number = 13, neutrons =n (unknown), atomic mass = 27. On
the board, substitute the values into the equation: 13 + n = 27. Then have a volunteer
solve for n. Students should easily be able to determine that n, the number of
• Review sample problem 2: Write a “skeleton” Bohr-Rutherford diagram on the board to fill in as you add information. Ask, How many protons in N? (7) How many
neutrons? (7) Write 7 p+ and 7 n0 in the nucleus of the diagram. How many electrons are there in all? (7) How many electrons are in the first orbit? (2) Draw the 2
electrons in the first (inner) orbit. How many electrons in the next orbit? (5) Draw the
5 electrons in the outer orbit. Make sure that electrons are paired, if possible.
• Provide additional practice by having all students complete the first page of BLM 6.7- 1: Practice Problems independently. Encourage students to complete the second page
to improve their understanding of isotopes and calculations of atomic mass. • Refer students to Communication Example 1: Write a “skeleton” Bohr-Rutherford
diagram. Ask, How many protons in F? (9) How many neutrons? (10) W Write 9 p+
and 10 n0 in in the nucleus of the diagram. How many electrons are there in all? (9) How many electrons are in the first orbit? (2) Draw the 2 electrons in the first (inner)
orbit. How many electrons in the next orbit? (7) Draw the 7 electrons in the outer
orbit.
• Have students complete Try This: Family Resemblances in the Periodic Table Try This: Family Resemblances in the Periodic Table
Skills: Performing, Analyzing, Communicating
Purpose: To determine patterns in Bohr-Rutherford diagrams for the first 20 elements of
the periodic table
Equipment and Materials (per student): periodic table, paper, pen or pencil Notes:
• If students draw their diagrams on index cards instead of directly on the paper, they
can use the cards in one of the Differentiated Instruction activities (below).
• Encourage students to check the accuracy of each drawing, making sure that the
number of protons in each atom equals the number of electrons.
Suggested Answers:
A. (i) Each element has a different total number of electrons, but there are eight electrons in the outermost orbit of each (except for He). Going down a family, the number of orbitals increases.
(ii) These elements all have one electron in their outermost orbit. (iii) These elements all have two electrons in their outermost orbit. (iv) These elements all have seven electrons in their outermost orbit.
B. (i) The alkali metals all have seven more spaces for electrons in their outermost orbit, whereas the noble gases have no more spaces for electrons in their outermost orbit. (ii) The halogens have one space left for electrons in their outermost orbit, whereas,
again, the noble gases have no more spaces for electrons in their outermost orbit. (iii) The alkaline earth metals all have one more electron in their outermost orbit
Extend and Assess
• To wrap up the lesson, review Table 1 on page 239 of the Student Book. Review the “theory” and “evidence” for each item and have students respond to each question in their notebooks. For example, you might ask, What is the evidence that each element is made of different kinds of atoms? Students should recognize that the fact that
different elements have very different properties is evidence that different elements have different kinds of atoms. Repeat with other examples, such as, What is the evidence that most of an atom is empty space? Students should respond that in
Rutherford’s gold foil experiment, most of the particles passed straight through the foil without hitting anything.
• To connect the material in this lesson with the next lesson, have students refer to their Bohr-Rutherford diagrams for a carbon atom. Tell them that carbon atoms can form different arrangements, creating several surprisingly different substances with remarkable properties. The four outer electrons in the atom are the basis for the different ways that carbon atoms chemically join with other carbon atoms and with atoms of other elements.
Check Your Learning Suggested Answers
1. Yes. The number of protons determines the chemical identity of an atomic particle. 2. No. Isotopes of a given element contain different numbers of neutrons.
3. (a) Correct
(b) Incorrect: The atomic mass must be greater than the atomic number, or close to equal to the atomic mass in the case of hydrogen.
(c) Correct, but only for the most common isotope of hydrogen, H-1. For the naturally occurring isotopes of all other elements, the mass number will exceed the atomic number due to the presence of neutrons.
(d) Incorrect: The number of protons may equal the number of neutrons in an atom. (e) Correct.
(f) Correct.
4. The most common isotope of Nickel, Ni-58, has 30 neutrons.
5. The first orbit can hold 2 electrons, and the second and third orbits can each hold 8 electrons. 6. Element name Element symbol Atomic number Mass number Number of protons Number of neutrons Number of electrons magnesium Mg 12 24 12 12 12 aluminum Al 13 27 13 14 13 phosphorus P 15 31 15 16 15 tin Sn 50 119 50 69 50 silver Ag 47 108 47 61 47
7. (a) (b)
(c) (d)
8. (a) False: The first four elements in the alkali metal family have different numbers of electron orbits.
(b) True (c) True (d) True
9. Elements in a given family have similar chemical properties, such as the alkali metals, which all react very vigorously with water. Each of the alkali metals has one electron in its outermost orbit and since the outermost electrons are the ones that participate in chemical changes, these elements react alike.
10. An element is a substance composed of atoms that all have the same number of protons in the atomic nucleus.
READING TIP