PERIODIC TABLE NOTES (from chapters 5 and 6) I. History of the Periodic Table As the number of elements began to grow, chemists needed a way to all of these elements. [In the 1700 s there were known elements. As new technologies were developed many new elements were discovered. In less than years the number of known elements had. A. J.W. Dobereiner: - Early - Grouped elements into sets of three s called. - He grouped elements like, Cl, Br, & I; Ca, Sr, & Ba. - [If you notice these groups on the current periodic table, they are all in the same row.] - The elements in the triad had similar properties. - The element placed in the middle has properties which are an of the other two elements. B. J.A.R. Newlands: - - There were known elements at this time. - He noticed that when the elements were arranged by increasing, the properties of the element were like those of the, the like the, 10 th like 3 rd, etc. - He called the pattern he saw the because the pattern repeated every element. Unfortunately his important recognition of the (or repetitive) pattern of element properties was not accepted because of its use of a term (octave-eight notes which make up a scale). C. Dmitri Mendeleev: D. Lothar Meyer: - - both of these scientists published nearly ways to classify the elements ** was given credit because he could demonstrate and the classification system. - Mendeleev eventually produced the periodic table, basically arranging the elements in order of. It was arranged so that elements in the same column had properties. His periodic table is very similar to the one we use today and because of this he is called the.
- Mendeleev sometimes the pattern of increasing atomic mass to keep the elements with properties in the same. - He switched sets of elements (one example is Co and Ni). [Note: these are out of order according to atomic mass on the periodic table.] He said that the atomic masses had been measured and when correctly measured, would match his table. - Mendeleev was also able to predict the and of missing elements. He correctly predicted the properties of (he called it Ekasilicon) which was later. E. The Periodic Law: Mendeleev was incorrect about the atomic masses being wrong. In 1913, did experiments from which he discovered the (the number of protons in the nucleus). He recognized that the correct way to arrange the elements was by atomic number (not as Mendeleev had done). Moseley then developed the which states that when elements are arranged in order of increasing, their chemical and physical properties show a pattern. II. Organizing the Periodic Table: [Because of the way that the periodic table is arranged, you can predict an elements properties by knowing its position on the table and what that position signifies.] The modern table has approximately squares. Each square includes certain information, like the, element, atomic, outer electron configuration, element name and other information depending on the table. The shape of the perio dic table comes from the. The two rows at the bottom actually fit into the table, but are placed at the bottom so the table can fit on one page and is not so. See Fig 5-8, pg 164 for a picture of the table with these two rows included.
A. Groups/Families - columns on the periodic table; - contain elements with similar - labeled with designation, or numbered 1-18. - We will use the A/B designation. [Please label your periodic table with the groups 1A 8A (or 0) and 1B 8B as it is done on the periodic table on the wall or on page 166-167 (red letters).] - A groups are referred to as the - B groups are called the - For the A groups, the number indicates the number of electron for that group. B. Periods - rows on the periodic table. - There are periods on the periodic table. -The two periods at the bottom are a part of periods (La-Vb) and (Ac-No). [Please label the periods on your periodic table as in done on the wall or pg. 166-167 in your book.] [There are a number of ways to group elements on the periodic table.] C. Metals/Nonmetals/Semimetals: [There is a division line on the periodic table called the stairstep. It is drawn in red on the wall table and starts on the box which contains B. Draw this on your own periodic table. This is what we use to identify metals and nonmetals.] Metals: Nonmetals: - located to the of the stairstep - are good conductors of and electricity - most are ( can be hammered into thin sheets) - most are ( can be drawn into fine wires) - have ( or shine) - metals are mostly at room temperature (except ) - are located to the of the stairstep. - are poor of heat and electricity - tend to be
- many are at room temperature - do not have luster ( or shine) Semimetals: - are located to (or touching) the stairstep. - have properties of both metals and non metals. Note: Al is actually a metal by its behavior, but we will call it a semimetal since it touches the stairstep. D. Group Names: Group 1A: Alkali Metals ( does not include H) - are metals, soft enough to be cut with a knife. - rapidly when exposed to air. - very reactive with, even water found in air; become more reactive as you move down the group. - stored under l to prevent reaction with moisture and oxygen - react with to form salts - never found in nature - These metals all have valence electron (outer energy level electron) - Readily lose one electron to form an ion with a charge - Outer electron configuration is followed by. Group 2A: Alkaline Earth Metals - Have all of the properties of. Are as similar as group 1A elements are. - Have valence electrons - Readily lose electrons to form an ion with a charge - Outer electron configuration is the followed by. - These elements are never found or uncombined in nature. Group 3A: Boron Group - Contains semimetals and metals - Have valence electrons - Readily lose electrons to form an ion with a charge - Outer electron configuration is the followed by
- Most important element in this group is Al. Group 4A: Carbon Group - Contains nonmetals, semimetals and metals - Have valence electrons - Will either valence electrons to become an ion with a charge, or valence electrons to become an ion with a charge. - Outer electron configuration is the followed by - Most important element in this group is C. Group 5A: Nitrogen Group - Contains nonmetals, semimetals and metals - Have valence electrons - Readily gain electrons to form an ion with a charge - Outer electron configuration is the followed by - Most important element in this group is N. Group 6A: Oxygen Group - Contains nonmetals, semimetals and metals - Have valence electrons - Readily gain electrons to form an ion with a charge - Outer electron configuration is the followed by - Most important element in this group is O. Group 7A: Halogens - Contains metals and semimetals - Have valence electrons - Readily gain electron to form an ion with a charge - Outer electron configuration is the followed by - All of these elements form as molecules (F 2, Cl 2, Br 2, I 2, etc). - These elements are very and therefore do not exist as elements in nature. - Some are very dangerous gases, like F and Cl. Group 8A ( or 0): Noble Gases - Are very, sometimes called gases. - Have valence electrons - Will not or electrons, because they already have a full set of 8 valence electrons.
- Outer electron configuration is the followed by Group B elements: Transition Metals - There is a great variety in this group. This group contains many of the common metals and metals used in. - They will lose a various number of electrons. - The outer configuration is basically. We will discuss this more later. Bottom Two Rows: Inner Transition Metals - Lanthanides: La-Yb - Actinides: Ac-No - There is a great variety in this group. - They will lose a various number of electrons. - The outer configuration is basically. We will discuss this more later. III. Outer Electron Configurations Valence electrons are electrons which occupy the principle. These electrons are basically responsible for an atoms behavior. As we have just discussed, atoms in the same have the same number of electrons. Because of this, all atoms within a certain group behave in a way. Since the valence electrons are the ones responsible for an atoms behavior, these are the electrons which are important for us to be aware of. In the last chapter we learned how to write the electron configuration for various atoms. This configuration showed us where each electron in the atom was located. Since we are really interested in the valence electrons, we can now write a electron configuration which will only indicate these valence electrons. This configuration is called the outer electron configuration. This outer electron configuration has 4 parts: 1. The which is determined by finding the noble gas (group 8A) in the period the period the atom you are working with is in. The noble gas is placed in brackets. See example below.
2. The is determined by the number if the outer electrons are in the s and p orbitals. If the electron is in the d orbital, the energy level is the. If the electron is in the f orbital, the energy is the. 3. The and 4. The in that orbital is determined from the group number. Recall: Group 1A: s 1 Group 2A: s 2 Group 3A: s 2 p 1 Group 4A: s 2 p 2 Group 5A: s 2 p 3 Group 6A: s 2 p 4 Group 7A: s 2 p 5 Group 8A: s 2 p 6 Transition Metals: s 2 d 1-10. To determine the number for the d orbital, you count over starting on the left of the B section. Inner Transition Metals: s 2 f 1-14.To determine the number for the f orbital, you count over starting on the left of the inner transition metal section. Example: Write the outer electron configuration for Na. 1. noble gas: 2. outer energy level: comes from the period #. 3 and 4. orbital and number of electrons: comes from group #. Now lets put it all together: Example: Write the outer electron configuration for I. 1. noble gas: 2. outer energy level: comes from the period #. 3. 3 and 4. orbital and number of electrons: comes from group #. Now lets put it all together:
**Note: you must put the energy level # in front of the s and p. Example: Write the outer electron configuration for Ni. 1. noble gas: 2. outer energy level: comes from the period #. 3 and 4. orbital and number of electrons: comes from group ( To determine the electron # for the d, we need to count over starting with Sc to Ni. Ni is the 8 th element in the B section, so it will be d 8. Remember the energy level number is also different for d (period # -1), so we will write 3 d 8 in our outer electron configuration.) Now lets put it all together: **Note: you must put the energy level # in front of the s and d. Example: Write the outer electron configuration for U. 1. noble gas: 2. outer energy level: comes from the period #. 3 and 4. orbital and number of electrons: comes from group. (To determine the electron # for the f, we need to count over starting with Th to U. U is the 3 rd element in the inner transition metal section, so it will be f 3. Remember the energy level number is also different for f (period # -2), so we will write 5 f 3 in our outer electron configuration.)
Now lets put it all together: **Note: you must put the energy level # in front of the s and d. You need to be able to write these outer electron configurations. To check your work, see the period table on pg.166-167 in your book. IV. Periodic Trends Many properties of elements will change in a predictable way as you move across and down the periodic table. These are called periodic trends. 1. Atomic Radius: is the distance from the center of an atoms nucleus to its outermost electron. As you move across a period from L to R, the atomic radius decreases. As you move down a group from top to bottom, the atomic radius increases. Remember that the largest atom is Fr and the smallest atom is F. Why: moving down a group, atoms have more electrons and more energy levels, so are bigger. Moving across a period, electrons are more strongly attached to the more positive nucleus, so atom is smaller. 2. Ionic Radius: is the distance from the center of an ions nucleus to its outermost electron. No convenient trend, but you should remember the following. When an atom loses electrons (to become a positive ion), the ion is smaller than the original atom. When an atom gains electrons ( to become a negative ion), the ion is larger than the original atom. 3. Ionization Energy: is the energy needed to remove one of an atom s electrons, or how strongly an atom holds its outermost electrons. As you move across a period from L to R, the ionization energy increases. As you move down a group from top to bottom, the ionization energy decreases. Remember that the atom with the highest ionization energy is F and the smallest ionization energy is Fr.
Why: Larger atoms(fr) hold electrons less tightly than small atoms (F), so the electrons in the larger atoms are easier to remove. 4. Electronegativity: reflects an atoms ability to attract electrons, or how much does an atom want. As you move across a period from L to R, the electronegativity increases. As you move down a group from top to bottom, the electronegativity decreases. Remember that the atom with the highest electronegativity is F and the smallest electronegativity is Fr. Why: Atoms with valence # s closer to eight want electrons more and therefore have a higher electronegativity.