Periodic Trends

  1.  What increases in equal steps of one from left to right in the periodic table for the elements lithium to neon?
    Atomic Number
  2. Which properties decrease down group 7 in the periodic table?
    • Electrongetivity
    • electron affinity
    • ionization energy
    • reactivity
  3. Which properties decrease down group 1 in the periodic table? 
    • Electrongetivity
    • electron affinity
    • ionization energy
    • melting point
  4. Valance Electrons in groups 1-8
    • 1:1
    • 2:2
    • 3A:3
    • 4A:4
    • 5A:5
    • 6A:6
    • 7A:7
    • 8A:8/0 He has 2 (exception)
  5. What is the trend across and down the  periodic table for atomic radius 
    Across– atomic radius decreases

    Down– atomic radius increases
  6. What is the trend across and down the periodic table for ionization energy
    Across– IE increases

    Down– IE decreases
  7. What is the trend across and down the periodic table for electronegativity
    Across– EN increases

    Down– EN decreases
  8. What two forces are responsible for the trends  AR, IE, EN?
    The two forces are the attractive force of the nucleus and the repulsive force of the electrons
  9. Give two reasons for increasing atomic radius down a group on the periodic table
    • Atomic radius increases going down a group because the electrons are added to energy levels that are farther from the nucleus and
    • Theouter electrons are shielded from the full attraction of the nucleus by the electrons in inner energy levels
  10. Give two reasons for lower ionization energy going down a group on the periodic table.      
    Ionization energy decreases going down a group because the electrons are farther from the nucleus b/c they are added to outer energy levels.

    Furthermore the electrons are shielded from the full attraction of the nucleus by the electrons in inner energy levels
  11. Give two reasons for the increase in ionization energy going across a period from
    left to right.
    • ionization energy increases across a period because the electrons are closer to the nucleus and because there are more protons attracting the electrons. 
    • The electrons are closer to the nucleus because within a period, they are added to the same energy level at the same rate that protons are added to the nucleus which increases the attraction of the electrons to the nucleus and drawing them closer.
  12. Give two reasons for the decrease in atomic radius going across a period from left to right. 
    the atomic radius decreases because nuclear charge increases and electrons within a period are added to the same energy level. 

    The increased nuclear charge pulls electrons closer – decreasing the radius.
  13. Explain why the second ionization energy for magnesium is lower than that for sodium. 
    • ionization energy generally increases across a period so this seems like an exception except that sodium only has one valence electron while magnesium has 2.  Sodium only needs to lose one electron to have a stable noble gas configuration and once it loses it, it is very unwilling to give up a second electron. 
    • Magnesium, on the other hand, easily gives up its second electron since it needs to  lose two to have the stable noble gas configuration.
    • Second ionization energy refers to the energy it takes to remove an electron after one electron as already been removed. 
  14. Explain why the first ionization energy for boron is slightly less than that of
    • Magnesium has two electron in the s sublevel while boron has two in the s sublevel and one in the p sublevel. 
    • Boron is more willing to give up its p than Mg is to give up one of its S electrons because there is some stability gained in having a filled sublevel.
    • B 1s2 2s2 2p1
    • Mg 1s2 2s2
  15. Explain why the first ionization energy of oxygen is slightly less than that of
    Ionization energy generally increases across a period so oxygen should have a higher ionization energy than nitrogen but it doesn’t.   If you look at the orbital diagrams for nitrogen and oxygen you will see that oxygen has one p orbital with two electrons and two other p orbitals with one electron each.  The added repulsion from the two electrons in one orbital and the fact that there is some stability gained by having a half-filled sub-level, explains why it is easier to take an electron from oxygen than nitrogen. Nitrogen already has a half-filled sub-level.
  16. Describe the change in atomic radius when anions form.
    • When anions form, electrons are gained.  The nucleus cannot overcome the added repulsion of the electrons and the radius increases.
    • The more negative, the larger the ion. 
  17. Describe the change in atomic radius when cations form.
    • When cations form, electrons are lost. With fewer electrons, there is less repulsion and the electrons are pulled closer to the nucleus. Therefore, for positive ions, radius decreases.
    • The more positive, the smaller the ion.
  18. Write reactions for lithium and water and potassium and water.
    Li+ H2O à LiOH + H2

    K+ H2O à KOH + H2
  19. State weather the resulting solutions will be acidic, neutral or alkaline.
    The solutions will be alkaline (basic) because of the presence of OH- ion.
  20. List some observations you could make while observing lithium react with water.
    • Fizzing (formation of hydrogen gas)
    • Heat and light (exothermic reaction)
    • Formation of alkaline solution
    • Metals seem to float on water and move around
  21. Reactivity with water increases/decreases going down the periodic table for group I.
  22. Explain the trend in reactivity in group 1 in terms of ionization energy.
    • Since ionization energy increases going  down a group it is easier for K to lose an electron than for Lithium to lose an electron, therefore K will be more reactive than lithium.  
    • **More energy is given off, though, in the lithium –water reaction because lithium has a higher charge density so more energy is released when water molecules surround the lithium atom.
  23. Suggest a reason that melting point decreases down group 1 on the periodic table. 
    • Group 1 elements are all metals and therefore bond via metallic bonding positive
    • nuclei in an electron sea.  Because of increased shielding the nuclei have less attraction for the electrons. 
    • Also, because there are more electrons, there are increased repulsion between the atoms. 
    • Cesium can be melted by holding it in your hand. 
  24. Suggest a reason that melting point increases down group 7 on the periodic table.
    • Group 7 are all non-metals and therefore bond via atomic network bonding. The electrons are not delocalized as they are in metals so the bonding is not as strong as it is in metals.  So this explains why fluorine and chlorine are gases at room temperature. As the number of electrons increase, there is increased dispersion forces which explainwhy bromine is a liquid and iodine is a solid. 
    • Dispersion forces are very weak, however, bromine is very volatile and iodine sublimes when open to the air.
  25. Explain the trend in reactivity in group 7 in terms of electron affinity
    • Electron affinity decreases going down a group because electrons are farther from the nucleus and are shielded by the inner electrons. Chlorine has a higher electron affinity than bromine which means that it would like electrons more than bromine. So, chlorine is likely to be more reactive than Br2 and I2. 
    • Chlorine is also a stronger oxidizing agent than bromine and bromine is a stronger oxidizing agent than iodine. 
  26. List some observations you could make when halide ions are mixed with silver ions. 
    • Most silver halide compounds are insoluble in water so a precipitate will form
    • AgCl is white
    • AgBr is cream/light yellow
    • AgI is yellow
    • AgF is soluble in water
  27. List the oxides formed by the period three elements.
    • Na2O
    • MgO
    • Al2O3
    • SiO2
    • P4O10
    • SO3
    • Cl2O7
  28. Describe the trend in acid/base properties from left to right
    The trend is strong base (Na2O) to weak base (MgO) to amphoteric (Al2O3) to weak acid (P4O10) to strong acid (SO3 + Cl2O7). These compounds become acidic/basic only when dissolved in water.
  29. Explain why SiO2 is considered an acid even though it is insoluble in water.
    • Water can’t break apart the strong Si-O bonds so it’s insoluble but is classified as an acid because when it reacts with a base it forms a neutral salt.  
    • Remember that the products of neutralization reactions are water and a salt.  So, if SiO2 reacts with a base and makes a neutral solution, it must be acidic
  30. Explain why Sodium Oxide forms a stronger base in water than Magnesium oxide.
    • Na2O separates into Na+ and O2-, O2- is a very strong base and takes a hydrogen from water. 
    • With Mg2+, the oxide is also strongly attracted to the more highly charged Mg2+ ion and therefore has less of an ability to steal protons from water.
  31. Define amphoteric and identify the periodic 3 oxide that is amphoteric.
    Amphoteric means a substance can act as an acid and a base depending on the conditions. Al2O3 is amphoteric.
  32. Likely to occur?
    Cl2+ I-
    yes Cl- + I2
  33. Likely to occur?
    Cl2+ Br-
    yes Cl- + Br2
  34. What is the shielding effect?
    Electrons between the nucleus and the valence electrons repel each other making the atom large
  35. How are Electronegativity and Electron Affinity different?
    Electronegativity attracts electrons in a chemical bond and electron affinity is energy change to form negative ions.
  36. Define/describe Electron Affinity
    • Preference for electrons
    • the energy change when an electron is added to the neutral atoms to form a negative ion
  37. Define/describe Electronegativity
    • The ability of an atom to attract electrons in a chemical bond.
    • In a group: increases UP because atoms with fewer energy levels can attract electrons  better.There is less shielding.
    • In a period: increases right because atoms can better attract electrons.
  38. Define/describe Ionization Energy
    • Energy required to remove an electron in order to make an ion.
    • Metals lose electrons more easily, they don't have a very strong E.N.C.
    • Nonmetals gain electrons to become stable because they have a higher E.N.C.
  39. Define/Describe Effective Nuclear Charge (E.N.C)?
    • How effectively the protons do or do not pull on their own electrons and the electrons of neighboring atoms.
    • IE, EN, EA, AR, and IR are a result of E.N.C
  40. Likely to occur?
    I2+ Br-
  41. Likely to occur?
    I2+ Cl
  42. Likely to occur?
    Br2+ Cl-
  43. Likely to occur?
    Br2+ I-
    yes Br- + I2
Card Set
Periodic Trends
Chemistry Test