Study Guide Chapter 7 PDF

Preview

Summary

Download Study Guide Chapter 7 PDF

Description

CHAPTER 7 Study Guide 1. The Periodic Table 2. Z effective ( Z eff) a. Each electron in an atom feels simultaneous attractions to the protons in the nucleus and repulsion by other electrons in the atom. For instance, the outermost electron in an atom feels attracted to the protons in the nucleus, but also repulsed by the other electrons in the inner shells that are between it and the nucleus. b. We say then, that the in-between electrons SHIELD the outermost electron from the nucleus. This makes it so that the outermost electron doesn’t ‘feel’ the protons in the nucleus as strongly as it otherwise would. c. Obviously inner electrons feel the pull from the nuclear protons more strongly than the outer shell electrons. The strength of attraction between an electron and the protons in the atom’s nucleus is called its effective nuclear charge, or Zeff. d. Can be calculated mathematically as follows: Zef = Z – S Where Z = number of protons in nucleus; and S = screening constant i. S = number of electrons found in orbitals that are lower in energy than the orbital of the electron in question. ii. So for valence electrons in Flourine 1. Step 1: Determine the number of protons in nucleus. This is Z. a. For Flourine, this is 9. 2. Step 2: Determine the complete electron configuration. a. For Flourine, this is: 1s22s2p5 3. Step 3: Describe which orbitals are the valence orbitals. Those are the outermost orbitals, which are the ones that have the highest n (principle quantum number). a. For Flourine, these ate the 2s and 2p orbitals 4. Step 4: Describe which orbitals are lower in energy than the valence orbitals. a. For Flourine, there is only one… the 1s orbital 5. Step 5: Count how many total electrons in the lower energy orbitals a. For fluorine, this is 2. And this is equal to S, the screening constant. 6. Step 6: Calculate Zef a. For fluorine, this is 9 – 2 = 7 iii. Can do the same for inner shell electrons with same equation. iv. Electrons in same orbitals as electron in question do not count. v. The larger the Zef, the more attraction to the protons in the nucleus. 3. Size of Atoms

a. Atoms get bigger as you go down a column in the Periodic Table due to more orbitals existing as you go down a column and principle quantum number, n, increases. b. Atoms get smaller as you go left to right across a row (period) because elements have more protons in their nucleus. These protons attract electrons more tightly drawing them closer to the nucleus making the atoms smaller.

4. Bonding Radii a. When two identical atoms bond, the distance between their nuclei, d, divided by 2, equals their individual ‘atomic radius’. b. Physicists have measured the individual atomic radii of diferent elements c. To measure the distance between the two bonding atoms (their bond length) we add up their individual atomic radii. 5. Isolelectronic Series a. Diferent ions, of diferent elements, can have the same number of electrons. i. For example, the following ions all have the same number of electrons. O-2, F-, Na+, Mg+2, Al+3 6. Electron Configuration of Ions a. Step 1: Derive electron configuration of neutral atom. b. Step 2: If cation; Starting at outermost orbital, remove same number of electrons as have been removed to form the ion. c. Step 3: If anion; Starting at the outermost orbital, add same number of electrons as have been gained to form the ion. 7. Ionization Energies a. It takes energy to remove electrons from an atom, even for atoms that want to give up their electrons. The energy required to remove an atom’s electron is called its ionization energy.

b. The first ionization energy is the energy required to remove the first electron from a neutral atom. c. The second ionization energy is the energy required to remove a second electron from a neutral atom. d. The larger the ionization energy, the harder it is to remove the electron. e. Electrons in higher energy orbitals are easier to remove, because they’re further away from the nucleus. i. For example, it’s easier to remove an electron from a 3s orbital than a 2s orbital.

8. Electron Affinities Ex. C(s)  C+(s) - eΔE = 1086 kJ/mol a. The positive ΔE means that energy must be put into the atom to remove its electron… Endothermic b. Most atoms can also gain electrons to become anions. The energy change in adding an electron to an atom is called the atom’s electron affinity. c. For most elements, energy is given of when an electron is added. i. For example, adding an electron to a sulfur atom gives of energy like this: S(s) + e-  S(s) ΔE = -200 kJ/mol 1. – ΔE signifies exothermic process d. Generally speaking, the more badly an element wants electrons, the more negative will be its electron affinity. 9. Metals, Nonmetals, and Metalloids a. Metals i. Shiny, malleable, ductile, and conduct heat and electricity ii. Elements show increasing metallic character as you go down and to the left of the Periodic Table.

iii. Have low 1st ionization energies. They easily lose electrons to become cations, means metals like to be oxidized. iv. Transition metals can often form ions with diferent charges. v. When metals bond to nonmetals they generally form ionic compounds vi. Most metal oxides are basic, meaning they react with water to form hydroxide salts. 1. Na2O (s) + H2O (l)  2 NaOH (aq) b. Non-metals i. Can be solid, liquid, or gas. They’re not lustrous and are generally poor conductors of heat and electricity. ii. When nonmetals bond with each other after they generally form molecular compounds iii. Because they have larger, negative electron affinities, they tend to gain electrons when they react with metals. iv. Most nonmetal oxides are acidic meaning they react with water to form acids like: CO2(g) + H2O (l)  H2CO3 (aq) c. Metalloids i. Have properties that are in-between those of metals and non-metals. 1. For example, Si looks like a metal but is not malleable. ii. Some are electron semi-conductors and are principle elements used in integrated circuits and computer chip....