CHAPTER 6 The Periodic Table

Transcription

1 CHAPTER 6 The Periodic Table 6.1 Organizing the Elements Mendeleev: listed the elements in order of increasing atomic mass and in vertical columns according to their properties. Left blank spaces for undiscovered elements Predicted the properties of undisovered elements. Henry Moseley: determined nuclear charge (atomic number) Arranged the periodic table by order of atomic number Significance of Subatomic Particles Protons: Define each atom Electrons: Determines the properties of atoms Neutrons: only alters the mass of an atom Periodic Organization 7 horizontal rows: periods Periods correspond to the principal energy level of atoms (Principal = Outer) Vertical columns: families/groups Corresponds to the number of outer energy level electrons 1

2 The Periodic Law When the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties 3 Classes of Elements Metals: good conductors of heat and electric current Nonmetals: poor conductors of heat and electric current Metalloids: properties of both. It depends on the conditions Metals About 80% of periodic table. High sheen and can reflect light. Most are ductile (make wires) Most are malleable and can be made into thin sheets. (Malleable = Bend) NonMetals Most are gases at room temperature. Bromine is a red liquid Solid nonmetals are brittle. Have a large variation of properties 2

3 6.2 Classifying the Elements 1A: Alkali metals 2A: Alkaline earth metals 7A: Halogens 8A: Noble Gases Configurations of the Groups Valence Electrons Electrons in s and p orbitals of outer energy level only. Show a pattern in the group. Increase across a period. Alkali Metals Group 1A s 1 =1 Valence electron Very reactive Rarely found alone (always bonding) Alkaline Earth Metals Group 2A Found in the crust of the earth. s 2 = 2 valence electrons 3

4 ELECTRON CONFIGURATIONS IN GROUPS There is a relationship between electron configuration of an element and its placement on the periodic table Blocks of Periodic Table Transition Metals Can bond with almost any non metal because it is willing to give up many electrons. d block Group B s, p, d, f blocks The location of elements within these blocks corresponds to the number of electrons in the outermost energy level F: 1s 2 2s 2 2p 5 Halogens (7A) Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 Group (6A) O: 1s 2 2s 2 2p 4 S: 1s 2 2s 2 2p 6 3s 2 3p 4 4

5 Alkali Metals (1A) Na: 1s 2 2s 2 2p 6 3s 1 (No 3p) Alkali Earth Metals (2A) Mg: 1s 2 2s 2 2p 6 3s 2 (No 3p) K: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 (No 4p) Ca: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 (No 4p) Noble Gases Group 8A Outermost s and p sublevels are completely filled (Inactive) Filled outer levels make atoms stable and inactive. s 2 p 6 = 8 valence electrons Valence Electrons Electrons in the last energy level of an atom. Include electrons in s and p orbitals only. Show a pattern in the group. Increase across a period. Representative Elements 1A-7A Outermost s or p sublevels are only partially filled (Group A elements) Represents elements from metals to non metals. Reactive Elements 8A not representative elements For any representative element, the group number is equal to the number of valence electrons. Groups 1A 7A. 5

6 1A: s 1 = 1 valence electrons 2A: s 2 = 2 valence electrons 3A: s 2 p 1 = 3 valence electrons 4A: s 2 p 2 = 4 valence electrons 5A: s 2 p 3 = 5 valence electrons 6A: s 2 p 4 = 6 valence electrons 7A: s 2 p 5 = 7 valence electrons 8A: s 2 p 6 = 8 valence electrons 1A 2A Group B Elements 3A 4A 5A 6A 7A 8A Transition Metals: (Group B) 1A 2A 3A 4A 5A 6A 7A 8A outermost s sublevel and the nearby d sublevel contain electrons All transition metals usually have the same number of valence electrons. Group B Elements Group B = Valence Electrons Do NOT determine properties Sc: [Ar] 4s 2 3d 1 Fe: [Ar] 4s 2 3d 6 Zr: [Kr] 5s 2 4d 2 Transition Metals 2 valence electrons Their properties are determined by the electrons in the d sublevel. Inner Transition metals outermost s sublevel and the nearby f sublevel generally contain electrons 6

7 Movement of Electrons 1A: 1 lost easily 2A: 2 lost easily 3A: 3 lost easily 5A: 3 gained 6A: 2 gained 7A: 1 gained Why do atoms form ions? Representative elements form ions to attain the electron configuration of a noble gas. Why do atoms form ions? Remember, as long as the proton numbers stays the same, the atom does not change. The electron number can change to attain the electron configuration of a noble gas. P = 10 N = 10 Neon vs. Sodium 1s 2 2s 2 2p 6 1s2 2s 2 2p 6 3s 1 P = 11 N = 12 Neon vs. Nitrogen 1s 2 2s 2 2p 6 1s 2 2s 2 2p 3 P = 10 N = 10 P = 7 N = 7 Ions for the groups Group 1A: Group 2A: Group 3A: Group 5A: Group 6A: Group 7A: 7

8 Charges of Atoms 1A: 1 + 2A: 2 + 3A: 3 + 5A: 3 6A: 2 7A: 1 BEFORE Na: 1s 2 2s 2 2p 6 3s 1 (Will lose 1 electron) B: 1s 2 2s 2 2p 1 (Will lose 3 electrons) P: 1s 2 2s 2 2p 6 3s 2 3p 3 (Will gain 3 electrons) F: 1s 2 2s 2 2p 5 (Will gain 1 electron) AFTER Na: 1s 2 2s 2 2p 6 ( Electron Configuration of Neon, Na + ) B: 1s 2 ( Electron Configuration of helium, B 3+ ) P: 1s 2 2s 2 2p 6 3s 2 3p 6 ( Electron Configuration of Argon, P 3 ) F: 1s 2 2s 2 2p 6 ( Electron Configuration of Neon, F ) Li Cl F # valence Gain or Lose # electrons Noble Gas Charge electrons electrons? gained/lost 6.3 PERIODIC TRENDS Nuclear Charge The properties of all atoms follow trends. These trends can be easily understood by using the periodic table. Trends go according to groups and periods. Proton Number Increases as atomic number increases. (Protons increase!) As proton number increases, the overall positive charge of the nucleus increases. 8

9 1A 2A Group B Elements 3A 4A 5A 6A 7A 8A Shielding Electrons Electrons between the nucleus and the valence electrons. They shield the valence electrons from the force of the nucleus. Na: 10 S: 10 Li: 2 N: 2 Shielding Electrons Shielding Electrons Total Electrons Valence Electrons Sodium: Total Electrons = 11 SE = Valence Electrons = 1 Shielding Electrons = 10 Sulfur: Total Electrons = 16 SE = Valence Electrons = 6 Shielding Electrons = 10 ENC = Total Shielding ENC = Total Shielding Shielding Electrons Total Electrons Valence Electrons Aluminum: Total Electrons = 13 SE = Valence Electrons = 3 ENC = Total Shielding Shielding Electrons = 10 Nitrogen: Total Electrons = 7 SE = Valence Electrons = 5 Shielding Electrons = 2 ENC = Total Shielding 1 e 8 e 2 e 1s 2 2s 2 2p 6 Total Electrons - Valence Electrons Na 11 Protons P = 10 N = 10 Shielding Electrons ENC= 9

10 1s 2 2s 2 2p 6 3s 1 Total Electrons - Valence Electrons 1s 2 2s 2 2p 6 3s 2 3p 3 Total Electrons - Valence Electrons P = 11 N = 12 Shielding Electrons P = 15 N = 16 Shielding Electrons ENC= ENC= 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 Potassium Arsenic Total Electrons = Total Electrons = Valence Electrons = Valence Eelctrons = Shielding Electrons= Shielding Electrons = ENC = ENC = Shielding Effect The further an energy level is from the nucleus, the more it is shielded from the attraction to the nucleus. The outer energy level is shielded by the inner energy levels. Atoms with more energy levels are larger. Effective Nuclear Charge Nuclear Charge Shielding Electrons The nuclear charge that the valence electrons actually feel. As the effective nuclear charge increases, the size of the atom decreases because the nucleus has more pulling power on the outer energy level of the atom. Effective Nuclear Charge Fluorine Atom: 9 protons & 9 electrons Valence electrons: 7 Shielding electrons: 2 Effective Nuclear Charge = +7 Nuclear Charge (9) Shielding Electrons (2) 10

11 Effective Nuclear Charge Carbon Atom: 6 protons & 6 electrons Valence electrons: 4 Shielding electrons: 2 Effective Nuclear Charge = +4 Nuclear Charge (6) Shielding Electrons (2) Nuclear Charge: Total Electrons: Valence Electrons: ENC vs. Energy Level Shielding Electrons: ENC: What if 2 atoms have the same ENC? Large atoms have more shielding electrons and are further away from nucleus. So, there is less pull on the valence electrons of a large atom than a small atom. Effective Nuclear Charge ENC = # Valence Electrons The greater the ENC, the more the nucleus pulls on the outer energy level. 11

12 Atomic Size/Radius Radius of an atom cannot be measured directly because you cannot see the outer electrons to measure the radius. Measured in picometers (10 12 ) Measuring Atoms Half the distance between the nuclei of two like atoms is the atomic radius Measure from nucleus to nucleus and take half Aluminum vs. Chlorine (Outer Energy Level Only) ENC +3 Both atoms have 3 energy levels, but chlorine's outer level shrinks due to the high ENC. ENC +7 Atomic Size/Radius (cont.) Increases to the left and down. More energy levels increases the size of the atom. Large ENC makes a small atom because there is more pull on outer energy level. Atomic Size/Radius (cont.) Atoms with more energy levels are larger, regardless of ENC. For example...iodine vs. Lithium Iodine Lithium ENC = +7 5 energy levels ENC = +1 2 energy levels 12

13 Atomic Size/Radius (cont.) Comparing 2 atoms in the same period High ENC atoms shrink because the the outer level is pulled very close to the nucleus. Low ENC atoms are large because the nucleus has little pulling power on the outer energy level. Group Trends Atomic size generally increases as you move down a group of the periodic table (More energy levels) Shielding increases for outer levels because there are more inner energy levels. Periodic Trends Atomic size decreases as you move from left to right across a period because the proton number increases. Shielding is constant because atoms do not gain energy levels across a period. (ENC increase) IONS An ion is an atom or group of atoms that has a positive or negative charge CATION: Positive Ions (Mg 2+ ) Lost one or more electrons ANION: Negative Ions (S 2 ) Gained one or more electrons Magnesium 1s 2 2s 2 2p 6 3s 2 Will lose 2 electrons. Mg 2+ Atom shrinks Protons vs. Electrons 12vs.10 Sulfur 1s 2 2s 2 2p 6 3s 2 3p 4 Will gain 2 electrons. S 2 Atom expands Protons vs. Electrons 16vs.18 TRENDS IN IONIC SIZE Cations are always smaller than their neutral form They can lose an energy level which decreases shielding. Electrons lose the tug of war with the protons. This is because the loss of outer shell electrons results in increased attraction by the nucleus for the fewer remaining electrons. Cations decrease across period 13

14 TRENDS IN IONIC SIZE Anions are always larger than the atoms from which they are formed. Electrons win the tug of war with the protons because the effective nuclear attraction for the nucleus is less when the number of electrons increases Anions decrease across period ANIONS Larger than neutral version. Negatively Charged Gain Electrons Nonmetals More they gain, bigger they get. CATIONS Smaller than neutral version. Positively Charged Lose Electrons Metals More they lose, smaller they get. Which is a smaller ion? Boron: 1s 2 2s 2 2p 1 Aluminum:1s 2 2s 2 2p 6 3s 2 3p 1 Which is larger? Oxygen: 1s 2 2s 2 2p 4 Oxygen Ion:1s 2 2s 2 2p 6 Ionic Charges Noble Gas Element Charge He: 1s 2 Ca: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Ne: 1s 2 2s 2 2p 6 N: 1s 2 2s 2 2p 3 Ar: 1s 2 2s 2 2p 6 3s 2 3p 6 Al: 1s 2 2s 2 2p 6 3s 2 3p 1 Kr: [Ar] 4s 2 4p 6 S: 1s 2 2s 2 2p 6 3s 2 3p 4 TRENDS IN IONIZATION ENERGY The energy that is required to overcome the attraction of the nuclear charge and remove an electron from an atom. Weak Pull = Low ENC = Low IE Strong Pull = High ENC = High IE 14

15 Ionization Energy(cont.) Increases right and up. Hard to remove electrons from atoms with outer energy level close to nucleus. Hard to remove electrons when the ENC is large. 1st IONIZATION ENERGY The energy required to remove the first electron from an atom. Depends on how many electrons exist in the outer energy level. See table 6.1 on page 173. TRENDS IN IONIZATION ENERGY 3 Things determine whether electrons can be removed ENC (effective nuclear charge) 2. The number of energy levels 3. The Tug of War after electrons have been removed 1st, 2nd, & 3rd Ionization Energy Na: 1s 2 2s 2 2p 6 3s 1 Remove 1 Remove 2 Remove 3 Mg: 1s2 2s2 2p6 3s 2 Remove 1 Remove 2 Remove 3 Al: 1s 2 2s 2 2p 6 3s 2 3p 1 Remove 1 Remove 2 Remove 3 2nd IONIZATION ENERGY Energy required to remove the second electron from an atom. Depends on how many electrons exist in the outer energy level. 15

16 2nd IONIZATION ENERGY It always requires more energy to remove the second electron than it does to remove the first. Element First Second Third Fourth Fifth Sixth Seventh Na 496 4,560 Mg 738 1,450 7,730 Al 577 1,816 2,881 11,600 P 1,060 1,890 2,905 4,950 6,270 21,200 S ,260 3,375 4,565 6,950 8,490 27,107 Cl 1,256 2,295 3,850 5,160 6,560 9,360 11,000 Ar 1,520 2,665 3,945 5,770 7,230 8,780 12,000 Group Trends Ionization energy decreases as you move down a group Li > Na > K Periodic Trends Ionization energy generally increases as you move from left to right across a period due to the number of protons increasing. TRENDS IN ELECTRONEGATIVITY The electronegativity of an element is the tendency for the atoms of the element to attract electrons when they are chemically combined with another element 16

17 Electronegativity.(cont.) Increases right and up. Closer outer energy levels attracts electrons more. Larger ENC attracts electrons more easily because the pull from the nucleus is greater. Each element is assigned an electronegativity value As you move across the periodic table the electronegativity increases Electronegativity decreases as you move down a group Period % Error = Predicted Accepted x 100 Accepted Group 7A Densities F: 1.7 Br: 3.12 At: 9.87 Cl = Factors determine most trends 1. ENC (Periodic Trends) Size Electronegativity Ionization Energy 2. Principal Energy Level (Group Trends) Size Shielding Electrons Attraction of valence electrons to the nucleus +1 Weak Pull +7 Strong Pull Density Which atom will attract the free electron? Metal Non Metal Which has the higher electron affinity? 17

18 Know these 7 Trends ENC Nuclear Charge Atomic Size (Atomic Radii) Shielding Electrons Ionization Energy Electronegativity Ionic Size What makes an atom large? What makes an atom small? What gives an atom a strong ENC? What gives an atom a weak ENC? What give an atom High Electronegativity? What gives an atom Low Electronegativity? Which atoms have High Ionization Energy? Which atoms have Low Ionization Energy? Possible Matching Terms Electronegativity Period Nuclear Charge Group Transition Metals Valence Electrons Inner Transition Metals Noble Gases Effective Nuclear Charge Halogens Ionization Energy Mendeleev Shielding Electrons p block Alkali Metals d block Alkali Earth Metals s block Periodic Law f block Atomic Size Group Trend Cation Mosely Anion Periodic Trend Summary of all trends is on Page 178. You must memorize all trends for the test! 18

19 Attachments key_configurations of the Groups.pdf