Chemistry Review Chapter 5

Chapter 5 Review 1. Identify two obvious trends going across the periodic table that you learned before we started chapter 5. What trend or trends can you identify going down a group that you learned before we started chapter 5. (Hint. Refer to Chapter 1, Section 1-3 of your book) a. Metals; Nonmetals Solids; Gases 2. Define periodic table. a. an arrangement of chemical elements based on their atomic numbers and similarity of properties 3. Who put the first successful periodic table together? Why did he call it “periodic”? a. Dmitri Mendeleev; He put the elements in recurring trends 4. In what order or by what trend did Mendeleev put the periodic table in? a. In the order of their atomic mass; Similarities in chemical properties 5. Which scientist corrected Mendeleev’s periodic table? What did he change? a. Henry Moseley (w/ help from Rutherford); arranged them by atomic number instead 6. Define periodic law. a. The statement that the physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of increasing atomic number. 7. What new group was added to the periodic table in the late 1800’s? a. Noble Gases 8. How is the modern periodic table arranged? (Hint: increasing/decreasing number?) Define this number. a. Increasing Atomic Number (The number of protons in the nucleus of an atom) 9. What is the common name for Group 1? Group 2? Group 7? Group 8? a. Alkali Family; Alkaline Family; Halogen Family; Noble Gases 10.What is the oxidation number for Groups 1-8? a. I(+1) II(+2) III(+3) IV(+4) V(-3) VI(-2) VII(-1) VIII(0) 11.The d-block is commonly known as what? What about the f-block? a. Transition Metals; Rare-Earth Metals 12.Why is Group 1 so highly reactive? Group VII? a. 1 Valence Electron and all the elements want it b. It only needs 1 to be happy and it’s desperate to get that 1 13.Why are noble gases largely unreactive? a. Because they are stable with 8 valence electrons and “happy” 14.What is the equation for ionization energy? What is the equation for electron affinity?

a. a + e a+ + e; a + e  a-+e 15.Define ionization energy. What is its trend on the periodic table? (period and group) a. he amount of energy that is needed to remove an electron from a molecule b. Period: Increases (left to right) Group: Decreases (top to bottom) 16.What is IE2? IE3? Which has more energy? a. The energies for removal of an additional electron from an atom; IE3 17.Define atomic radius. What is its trend on the periodic table? (period and group) a. half the distance between the atomic nuclei in a molecule consisting of identical atoms b. Period: Decreases (left to right) Group: Increases (Top to Bottom) 18.What are valence electrons? Where are they located? a. An electron in the outermost energy level of an atom; they are capable of forming bonds with other atoms 19.Why does the atomic radius increase when it is followed down a group? a. Because the number of energy levels increases therefore the farther the outermost valence electron the larger the atomic radius 20.Does Group 1 tend to lose or gain its valence electrons? Why? a. Lose; because it is easier to lose a valence electron then gain 7 21.Does Group 7 tend to lose or gain its valence electrons? Why? a. Gain; because it is easier to gain a valence electron then lose 7 22.Define electronegativity. What is its trend on the periodic table? (period and group) a. is the ability of an atom to attract toward itself the electrons in a chemical bond b. Period: Increases (Left to right) Group: Decreases (Top to Bottom) 23.Define electron affinity. What is its trend on the periodic table? (period and group) a. the energy needed to remove an electron from a negative ion to form a neutral atom or molecule b. Period: Increases (Left to right) Group: Stays constant, does not change 24.In ionic radius, what happens to the size of an atom when an electron is gained? When an electron is lost? a. Cation: Ionic Radius Decreases; Anion: Ionic Radius Increases 25.Distinguish between a cation and anion.

a. Cation – positive ion b. Anion – a negation ion

1. Discuss the importance of Mendeleev’s periodic law. 2. Identify each element as a metal, metalloid, or nonmetal. a. Fluorine - Nonmetal b. Germanium- Metalloid c. Zinc - Metal d. Phosphorous - Nonmetal e. Lithium - Metal 3. Give two examples of elements for each category. a. Noble gases - Krypton; Neon b. Halogens - Chlorine; Fluorine c. Alkali metals - Lithium; Sodium d. Alkaline metals - Strontium; Calcium 4. What trend in atomic radius do you see as you go down a group/family on the periodic table? What causes this trend? a. 5. What trend in atomic radius do you see as you go across a period/row on the periodic table? What causes this trend? a. 6. Circle the atom in each pair that has the largest atomic radius a. Al B b. S O c. Br Cl d. Na Al e. O F f. Mg Ca 7. Define Ionization Energy. a. he amount of energy that is needed to remove an electron from a molecule 8. Is it easier to form a positive ion with an element that has a high ionization energy or an element that has a low ionization energy? Explain. a. 9. Use of the concept of ionization energy to explain why sodium form a 1 +ion (Na+) but magnesium forms a 2+ ion (Mg2+). a. 10.What trend in ionization energy do you see as you go down a group/family on the periodic table? What causes this trend? a. 11.What trend in ionization energy do you see as you go across a period/row on the periodic table? What causes this trend? a. 12.Circle the atom in each pair that has the greater ionization energy. a. Li Be b. Na K c. Cl Si d. Ca Ba e. O Ar f. Li K 13.Define electronegativity. a. the attraction of an atom for electrons in a covalent bond

14.What trend in electronegativity do you see as you go down a group/family on the periodic table? What causes this trend? a. 15.What trend in electronegativity do you see as you go across a period/row on the periodic table? What causes this trend? a. 16.Circle the atom in each pair that has the greater ionization energy. a. Ca Ga b. Li O c. Cl S d. Br As e. Ba Sr f. O S 17.Define Electron Affinity. a. the energy needed to remove an electron from a negative ion to form a neutral atom or molecule 18.What trend in electron affinity do you see as you go down a group/family on the periodic table? What causes this trend? a. 19.What trend in electron affinity do you see as you go across a period/row on the periodic table? What causes this trend? a.