Honors Chemistry – Periodic Table PDF

Honors Chemistry – Periodic Table Tom Lehrer's 'The Elements' animated.mp4 https://www.youtube.com/watch?v=zGM-wSKFBpo The NEW Periodic Table Song (Updated).mp4 https://www.youtube.com/watch?v...

Honors Chemistry – Periodic Table Tom Lehrer's 'The Elements' animated.mp4 https://www.youtube.com/watch?v=zGM-wSKFBpo The NEW Periodic Table Song (Updated).mp4 https://www.youtube.com/watch?v=VgVQKCcfwnU History: Mendeleev -arranged the Periodic Table by atomic mass Moseley (current table) -The modern Periodic Table is arranged by increasing atomic number Reminder: atomic number is the number of protons, only! ______________________________________________________________________________ Basic Setup: -Metals are located to the left of the step -Nonmetals are located to the right of the step -Semimetals/metalloids are elements that border the step Properties of Metals: Properties of Nonmetals: -amount: more than half of the elements -location: to the right of the step are metals -number of valence electrons: -location: to the left of the step relatively high number of NOTE: Hydrogen is NOT a valence electrons (5-8) metal (even though it is written in the -atoms tend to gain electrons to get upper left of the table) -number of valence electrons: relatively a full valence shell and form low number of valence negative ions, anions electrons (1-3) -type of ions formed: negative, anions -can bond with other NM -atoms tend to lose electrons to -poor conductors; good insulators form positive ions called cations -states of matter: (You must know which NM are solid, which are liquid, and which -type of ions formed: positive, cations are gases) -states of matter: all solid except -gas: Hydrogen, mercury Nitrogen, Oxygen, Fluorine, -the only liquid metal at room temp: Chlorine, Helium, Neon, Mercury = Hg Argon, Krypton, Xenon, -metals are hard because: the particles are packed so closely together Radon -metals are excellent -liquid: Bromine is the only conductors because: they have liquid NM -Solid: carbon, freely moving electrons (excellent conductors of heat and phosphorus, sulfur, electricity) selenium, iodine, astatine Malleable, ductile, high -solids NM are brittle and soft melting points, luster Reminder: Which NM are naturally diatomic? 2 Properties of Semimetals/Metalloids: -have properties of both M and NM -you must be able to identify all semimetals by name and symbol: (B) Boron, (Si) Silicon, (Ge) Germanium, (As) Arsenic, (Sb) Antimony, (Te) Tellurium Aluminum: looks like it would be a metalloid by its location, but it is NOT. Aluminum is a metal! Aluminum is NOT a semimetal even though it borders the staircase. Groups: -columns (top to bottom) -1-18 -aka families b/c: members (elements) of the same family have similar characteristics -members (elements) of the same group have the same number of: valence e- -elements of the same group react similarly b/c: they have the same # of valence e- SAME GROUP = SAME # VALENCE e- = SIMILAR PROPERTIES -Group 1 (I A) family name: Alkali Metals So reactive that: these metals are naturally only found in compounds 1 valence electron (s1) Lose 1 valence electron to become a cation with a +1 oxidation state Stored in oil so it doesn’t react (explode) upon contact with moisture in the air These metals react with water to form H2(g) and a basic solution (OH-) Li(s) + H2O → Li(OH)2 + H2 3 -Group 2 (II A) family name: Alkaline Earth Metals Not as reactive as group 1 metals but also very reactive that they exist naturally in compounds 2 valence electrons (s2) Lose 2 valence electrons to become a cation with a +2 oxidation state -Groups 3-11: Transition Metals (incomplete d sublevels) Produce colors when dissolved in water (green, blue, yellow, etc.) s and d sublevels are very close together. The electrons are in constant transition! Valence electrons can be lost from s and/or d orbitals which is why they have multiple positive oxidation states -Group 12: Heavy Metals Zn, Cd, Hg -Group 17 (VII A): Halogens (s, l, and g) So reactive that they naturally exist only in compounds 7 valence electrons (s2p5) Gain an electron to become an anion with a -1 oxidation state Some have multiple (+ and -) oxidation states as a result of bonding in different ways Fluorine – pale yellow greenish gas Chlorine – yellow gas Bromine – dark brown-ish liquid turns to dark yellow/brown gas Iodine – silver-ish solid turns into purple gas 4  -Group 18 (VIII A or 0): Noble gases/Inert gases don't react/are inert because they have full valence shells (s2p6) or (s2 for He) **Kr and Xe they have been found to react with Fluorine, not common but has happened!! -Lanthanide and Actinide series Bottom of the period table f sublevels being filled MOST ACTIVE (REACTIVE) METAL IS FOUND IN THE LOWER LEFT CORNER OF THE TABLE (Fr). MOST ACTIVE (REACTIVE) NONMETAL IS FOUND IN THE UPPER RIGHT (F) (remember: noble gases do not react) Periods: -rows (left to right) -represent: the # of occupied p.e.l. (shells) -properties: change across a period from M to SM to NM Reminders: -an atom will never have more than 8 valence electrons -ground state: when all electrons are in their lowest possible energy level. Electron Dot Diagrams/Lewis Dot Diagrams: -2 parts to a LDD: 1: kernel: represents everything except the valence e- 2: dots: represent valence e- 5 Electron Dot Diagrams for atoms in the ground state: -charge: atoms are neutral! So, the e- dot diagrams for atoms in ground state will never have a charge written (charge is 0). -brackets ([ ]): e- dot diagrams for atoms in the ground state will never have brackets. Brackets indicate that e- have been taken in. This is not the case for atoms in the g.s. -dots: represent the number of valence electrons. Valence e- are the e- in the outermost shell. Examples: Draw the Lewis dot diagram for the following atoms in the ground state. Sodium Magnesium Aluminum Silicon Phosphorus Sulfur Chlorine Argon Electron Dot Diagrams for metallic ions (positively charged ions): -positive ions -cations Dots: no dots Brackets: no brackets Charge: written to the upper right of the symbol. Examples: Draw the Lewis dot diagram for a: Sodium ion Magnesium ion Aluminum ion Electron Dot Diagrams for nonmetallic ions (negatively charged ions): -negative ions -anions Dots: 8 dots Brackets: always. This shows e- have been taken in. Charge: written to the upper right of the symbol OUTSIDE the brackets. Examples: Draw the Lewis dot diagram for a: Phosphide ion Sulfide ion Chloride ion 6 Ionic Radii in comparison to the corresponding atomic radii: Metals: -form positive ions -The atomic radius of a metal is larger than the ionic radius of the same element because the atom of a metal has one more energy level than the ion of the same element. (The ion has one less energy level than the atom of the corresponding element). Example: sodium Nonmetals: - form negative ions -The atomic radius of a nonmetal is smaller than the ionic radius of the same element because the atom of a nonmetal has fewer electrons than the ion of the same element. As a result of having fewer electrons, the nucleus has a stronger pull on the valence shell which in turn creates a smaller radius as compared to the radius of the ion for the same element. (The ion of a nonmetal has a larger radius than its corresponding atom because the ion has more electrons. As a result of having more electrons the nucleus has a weaker pull on the valence shell which in turn creates a larger radius as compared to the radius of the atom for the same element. Example: fluorine 7 Reminders on Ions: -a charged particle -the only difference between an atom and an ion of a given element is the number of (valence) electrons. -M tend to lose their valence e- to become positive ions. This means metal atoms lose an energy level. So, metal atoms will have a larger radius than the ion of the same element. -NM tend to gain e- to become negative ions. This means NM ions will have a larger radius than the atom of the same element. Ion electron configuration: 1. Write the electron configuration for the sodium ion. 2. Write the electron configuration for a fluoride ion. 3. Write the electron configuration for S2- 4. Write the electron configuration for Ca2+ Determining number of protons and electrons 1. S2- 2. Ca2+ 3. Al3+ 4. N3- 5. Phosphide ion different forms of the same element. This means different Allotropes: properties and different structures. The only thing the same is the same element. Ex: Carbon: diamond, graphite (conducts), coal Oxygen: O2 and O3 Phosphorus: white, red, and black Periodic Law: properties of the elements are periodic functions of their atomic # (properties are in a pattern) https://www.youtube.com/watch?v=1PSzSTilu_s&list=PL65159266CFC74682&index=15 Mark Rosengarten’s “Elemental Funkiness” 8 Trends 1. Atomic Radius -the radius is determined as 1/2 the distance between two nuclei. -shells represent distances from the nucleus. -more p.e.l. = valence e- farther from the nucleus than an atom with fewer p.e.l. -more shells = larger radius -As you go from top to bottom down a group, atomic radius increases b/c more e- shells (pel) are present. As the # of shells increases, the inner (non-valence) shells shield the valence shell from the nucleus. So, the nucleus has a weaker pull on the valence shell. (this has nothing to do with the number of valence e-) -the stronger the pull from the nucleus, the smaller the radius. Li Na K Rb Cs ____________________________________________________________________________ -As you go from L to R across a period, the number of shells remains the same. -As you go from L to R across a period, the nucleus becomes stronger (more protons). -The stronger the pull from the nucleus, the smaller the radius. -So, as you go L to R across a period, atomic radius decreases Li Be B C N O F First: more e- shells = weaker attraction between nucleus and valence e- = larger radius IF AND ONLY IF # e- SHELLS ARE EQUAL, then… Then: more protons = stronger attraction between nucleus and valence e- = smaller radius 2. Reactivity a) Metals (metallic character) (REMINDER: Hydrogen (H) is NOT A METAL!) -as you go from L to R (increasing atomic #) across a period, metallic character decreases -as you go from Top to Bottom (increasing atomic #) down a group, metallic character increases 9 Reactions of alkali metals with water https://www.youtube.com/watch?v=m55kgyApYrY b) Nonmetals (nonmetallic character) -as you go L to R across a period, NM character increases -as you go T to B down a group, NM character decreases 3. Ionization Energy (I.E.) -is the amount of energy needed to remove the most loosely held electron (1st ionization energy, 2nd ionization energy, …) REMINDER: Metals want to lose e-. Therefore, metals will not require a lot of energy to remove one. So, metals will have low I.E. REMINDER: Nonmetals want to gain e- (NOT lose e-). Therefore, nonmetals will require a lot of energy to remove an e-. So, nonmetals will have high I.E. -As you go from top to bottom ↓(incr. atomic #) down a group, I.E. decreases (see group 15) because the atomic radius is larger and therefore the nucleus has a weaker hold on the valence shell. So, less energy is needed to remove an electron. 10 -As you go from L to R →(incr. atomic #) across a period (M --> SM --> NM), ionization energy GENERALLY increases, but there is a slight change between s and p sublevel filling: Na Mg Al Si P S Cl Ar 496 738 578 787 1012 1000 1251 1521 Lower IE = less energy needed to remove an e- = more reactive metal. First Ionization Energy compared to Second Ionization Energy: Na Mg Al st (1 IE=496kJ) st (1 IE=738kJ) 1st IE = 578 kJ Na + 496 kJ → Na+ Mg + 738 kJ → Mg+ 2nd IE= 1,820 kJ 3rd IE= 2,750 kJ (2nd IE=4,560kJ) (2nd IE=1,450kJ) 4th IE=11,600 kJ Na+ + 4,560 kJ → Na2+ Mg+ + 1,450 kJ → Mg+2 (3rd IE=7,730 kJ) Mg2+ + 7730 kJ → Mg+3 Why is the 2nd IE SO much Why is the 3rd IE SO much Why is the 4th IE SO much higher than the 1st IE for higher than the 2nd IE for higher than the 3rd IE for group 1 metals? group 2 metals? group 3 elements? 11 4. Electronegativity - is the affinity for e-/tendency to take in e-/ability to attract an e-/"love of e-" -the higher the electronegativity, the more likely the atom will be able to take in an e-. -Electronegativity is an arbitrary scale (0-4) -Fluorine (F) is the most reactive NM. *- Fluorine has the highest electronegativity = 4.0 So, NM will have higher electronegativity values than M. -As you go →across a period, electronegativity increases -As you go↓down a group, electronegativity decreases Ionization energy and electronegativity follow the same trend! If an atom wants to take in an e- (high electronegativity) the atom will require a lot of energy to remove an e- (IE) The trend for atomic radius is opposite to that of IE and electronegativity. Comparing radius of different species: Using your knowledge of attraction between nucleus and valence e-, list the following in expected ascending (increasing) order of radius. O _____ _____ _____ _____ _____ _____ _____ _____ _____ O2- smallest largest F F- hint: write out electron config and atomic # Ne Na Na+ Mg Mg2+ 12