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Chapter 6 - The Periodic Table
 

Chapter 6 - The Periodic Table

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This presentations covers all of Chapter 6 and the organization of the Periodic Table of the elements.

This presentations covers all of Chapter 6 and the organization of the Periodic Table of the elements.

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    Chapter 6 - The Periodic Table Chapter 6 - The Periodic Table Presentation Transcript

    • Chapter 6: The Periodic Table By Kendon Smith Columbia Central HS Brooklyn, MI
    • A. Development of the Periodic Table1. The first periodic table: (p. 156) a. Created by Dimitri Mendeleev in the 1800’s. b. Elements were listed in order of increasing atomic mass. c. Elements with similar properties were listed side by side and it was discovered that there were repeating patterns to these properties. d. Mendeleev left blank spaces where the pattern showed elements should exist. These elements were later discovered!
    • A. Development of the Periodic Table2. The modern periodic table: a. Rearranged by Henry Mosely in 1913. b. Elements were listed in order of increasing atomic number because the atomic masses didn’t always fit the pattern.
    • B. Organization of the Periodic Table1. The periodic table is arranged into horizontal rows called periods and vertical columns called groups or families.2. Periods: a. Periods contain different numbers of elements. Period 1: 2 Period 2: 8 Period 4: 18 Period 6: 32 b. Properties of elements change as you go across a period. c. The pattern of properties within a period repeats in each period. This is called the periodic law. d. Each period represents the number of energy levels holding the electrons of those elements.
    • B. Organization of the Periodic Table3. Groups: a. Elements in groups have similar chemical and physical properties. b. Each group is identified by a Roman numeral and letter A or B. We usually just number them from 1 to 18.Ex. Group IIA = 2 Group IIIA = 13 Group VIIA = 17
    • C. Metals, Nonmetals, and Metalloids - All elements are classified into 3 main categories based on their properties.1. Metals a. Metals exhibit the most metallic properties. - Good conductors of heat and electricity. - Surface can be polished to have a high luster or sheen. - Most are solids at room temperature. (Except: Mercury) - Many are ductile and can be drawn into wires. - Most are malleable and can be pounded into thin sheets.
    • C. Metals, Nonmetals, and Metalloids2. Nonmetals a. Nonmetals tend have properties that are opposite to those of metals. - Most are gases at room temperature. Bromine is a liquid. Sulfur and phosphorus are solids. - Most are poor conductors of heat and electricity. (Except: Carbon) - Most tend to be brittle and will shatter if it with a hammer. b. Nonmetals form the basis of organic chemistry because they make up so many compounds in living things.
    • C. Metals, Nonmetals, and Metalloids3. Metalloids a. Metalloids, called semi-metallic elements, are between the metals & nonmetals. b. Under certain conditions, metalloids may behave like metals or nonmetals. - Metalloids have mixed properties.
    • D. Important Periodic Table Groups and Properties: 1. Alkali Metals: Li, Na, K, Rb, Cs, Fr a. Alkali metals are found in Group 1A or 1. - All alkali metals have one valence electron. b. These metals are highly reactive and never found free in nature. c. All alkali metals react with water to release hydrogen gas. d. Li, Na, and K have densities less than 1.0 and float on water.
    • D. Important Periodic Table Groups and Properties: 2. Alkaline Earth Metals: Be, Mg, Ca, Sr, Ba, Ra a. Alkaline-earth metals are found in Group 2A or 2. - All alkaline earth metals have two valence e-. b. These metals are very reactive, but not as much as the alkali metals.
    • D. Important Periodic Table Groups and Properties: 3. Transition Metals – Fe, Ag, Cu, Au, Ti, Cr… a. Transition metals are found in Groups 3B – 2B or 3 - 12. - Characterized by having electrons in their d orbitals. - Most transition metals have two valence electrons, a few only have one.Examples: [Ar]4s23d6 = iron [Ar]4s13d5 = chromium (exception) *Note: The idea of valence electrons is not very useful with transition metals!
    • D. Important Periodic Table Groups and Properties: 4. Inner Transition Metals – U, Pr, Nd, Eu, Yb, Cf… a. Also called the Rare Earth Metals b. Inner transition metals are usually shown in 2 rows under the main table. c. Characterized by having electrons in their f orbitals. d. Elements beyond uranium are all man made and radioactive.
    • 5. Halogens: F, Cl, Br, I a. These elements are found in Group 7A or 17. - Characterized by having 7 valence electrons. b. Halogens are the most reactive of the non-metals. c. Halogens react vigorously with many metals to form a salt compound called a metal halide. - A common metal halide salt is NaCl – sodium chloride Reactions: 2 Na + Cl2 → 2 NaCl NaOH + HCl → NaCl + H2O
    • 6. Noble Gases: He, Ne, Ar, Kr, Xe, Rn a. These elements are found in Group 8A or 18. - Characterized by having 8 valence electrons. b. Also called the inert gases - means “inactive” c. These elements are the least reactive of all the elements.7. Representative Elements: - This name is sometimes given to elements in groups 1A – 7A, where the group number equals the number of valence electrons.
    • E. Periodic Trends1. Atomic Size: a. Usually expressed as an atomic radius measured in picometers. b. Atomic radius is equal to half the distance between the nuclei of two atoms of the same element when they are joined.
    • c. Atomic size increases as you move down a group. - Going down a row adds another energy level of electrons. - Lower energy levels shield the outer level from the positive charge pulling it towards the nucleus.Ex: Chlorine – 102 pm Bromine – 120 pm + +Lower levels shield the outer levels!
    • d. Atomic size decreases as you move from left to right across a period. - Atoms are increasing in mass and atomic number, but get smaller! - Protons are being added to the nucleus, making it more positive! - Electrons are being added to the same principal energy level! * There is no change in shielding effect of the inner levels. - Increased nuclear charge pulls valence electrons closer to the nucleus.
    • Ex: Phosphorus – 109 pm Chlorine – 102 pm 15 protons! 17 protons! + +More protons in the nucleus makes the positivecharge “stronger” and it pulls the electrons in a bittighter. More electrons, but closer to nucleus!
    • 2. Ions:a. Ions are atoms or groups of atoms that have a positive or negative charge. - Neutral = same number of protons and electrons - Positive atom = more protons than electrons - Negative atom = more electrons than protonsb. Ions can only form when atoms gain or lose electrons. - Gaining electrons makes an atom negative. - Losing electrons makes an atom positive.
    • c. Cations: - Cations are positively charged ions. - Metals tend to form positive ions by losing electrons. - Cation charges are written as a number with a plus sign. Examples: Sodium metal loses 1 e- Calcium metal loses 2 e- to become the sodium ion. to become the calcium ion. Na1+ Na+ Ca2+
    • c. Anions: - Anions are negatively charged ions. - Nonmetals tend to form negative ions by gaining electrons. - Anion charges are written as a number with a minus sign. Fix this in your notes! Examples: Chlorine gains 1 e- to Phosphorus gains 3 e- to become the chloride ion. become the phosphide ion. Cl1- Cl- P3-
    • 3. Ionization Energy: a. Ionization energy is the energy required to remove an electron from an atom. b. 1st ionization energy = remove 1st e- = 1+ charge c. 2nd ionization energy = remove 2nd e- = 2+ charge d. 3rd ionization energy = remove 3rd e- = 3+ charge - Higher ionization energy means electrons are harder to remove! - Harder to remove = less likely to form that ion!
    • Element 1st Ionization 2nd Ionization 3rd Ionization Na 496 kJ/mol 4565 kJ/mol 6912 kJ/mol Mg 738 kJ/mol 1450 kJ/mol 7732 kJ/mol K 419 kJ/mol 3069 kJ/mol 4600 kJ/mola. How many electrons will Na most likely lose? 1 e- What ion will that make? Na1+b. How many electrons will Mg most likely lose? 2 e- What ion will that make? Mg2+c. How many electrons will K most likely lose? 1 e- What ion will that make? K1+
    • e. Periodic Trend: First ionization energy decreases from top to bottom in a group and increases from left to right across a period.2 Ways to Think About Ionization Energy: - Atoms with fewer electrons hold on to the more tightly. (As they lose them it goes up!) - As atoms get larger, valence electrons are farther from the nucleus and easier to pick off.
    • 4. Ionic Size: (This one is easy!)a. As atoms lose electrons to become cations, their electron clouds get smaller! - Nucleus pulls harder on the remaining electrons.b. As atoms gain electrons to become anions, their electron clouds get bigger! - Nucleus can’t pull as hard on extra electrons.Sodium → Sodium ion Chlorine → Chloride ion
    • e. Periodic Trend: Ionic size increases from top to bottom in a group and ionic size of cations and anions decreases across a period.* Note: Anions are always larger than cations in the same period!Example: Which ion has a larger size? a. Sodium ion or Magnesium ion? Na1+ Mg2+ b. Magnesium ion or Phosphide ion? Mg2+ P-3 c. Sodium ion or Potassium ion? Na1+ K1+
    • 5. Electronegativity:a. Electronegativity is the ability of an atom to attract electrons. Higher EN = nucleus has a stronger pull on electronsb. Values range from 0.7 (Francium) to 4.0 (Fluorine)c. Periodic Trend: Values tend to decrease from top to bottom in a group and increase from left to right across a period. - Metals easily lose electrons = low EN values. - Nonmetals usually gain electrons = high EN values
    • Partial Table of Electronegativity Values Note: Why do Noble Gases not have EN values?
    • Atomic radius decreases Ionization energy & Electronegativity IncreaseAtomic RadiusincreasesIonization EnergydecreasesIonic SizeincreasesElectronegativitydecreases
    • THE ENDReady for your test soon?