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Chemistry Unit 4 Review Know these things: 1. How to write the electron configuration for exceptional elements. 2. Write the electron configuration for ions. 3. Relate an electron configuration to an element’s position on the periodic table. 4. Find the number of valence electrons for an element. 5. The definitions and trends for reactivity, atomic radius, ionization energy, electron affinity, and electronegativity. A number of physical and chemical properties of elements can be predicted from their position in the Periodic Table. Among these properties are Ionization Energy, Electron Affinity and Atomic/ Ionic Radii. These properties all involve the outer shell (valence) electrons as well as the inner shell (shielding) electrons. Electrons are held in the atom by their electrostatic attraction to the positively charged protons, the nuclear charge, Z. However, not all electrons in an atom experience the same nuclear charge. Those closest to the nucleus experience the full nuclear charge and are held most strongly. As the number of electrons between the nucleus and the valence electrons increases, the apparent nuclear charge decreases, due to the"shielding" of these inner shell electrons. Going down a group increases the value of n, and increases the number of inner shell electrons. This leads to better shielding and a weaker attraction between the nucleus and the outer shell electrons. Going across a period leads to a larger nuclear charge, as the number of protons increases. There is also an increase in the number of valence electrons, but electrons in the same shell are poor at shielding each other. Going across a row generally leads to a stronger interaction between the nucleus and the valence electrons. The type of orbital holding the shielding electrons is also important. The s orbitals are said to be penetrating - they have electron density close to the nucleus. The best shielding comes from s orbitals, followed by p, d and f. Ionization Energy Ionization energy is the energy required to remove an electron from a gaseous atom in its ground state. This is related to how "tightly" the electron is held by the nucleus. The higher the ionization energy, the more difficult it is to remove the electron. For a many-electron atom, the energy required for the reaction: Energy + X (g) → X+ (g) + e- First Ionization Energy (I1) Since this requires an input of energy, it is an endothermic reaction, with a positive energy value. The energy required for the reactions: Energy + X+ (g) → X 2+ (g) + eEnergy + X 2+ (g) → X 3+ (g) + e- Second Ionization Energy (I2) Third Ionization Energy (I3) When an electron is removed from an atom the repulsion between the remaining electrons decreases. The nuclear charge remains constant, so more energy is required to remove another electron from the positively charged ion. This means that, I1 < I2 < I3 < ..., for any given atom. Going down a group the electrons become increasingly easy to remove, since they are at an increasing distance from the nucleus, with increasing numbers of shielding inner electrons. So, the ionization energies decrease. Going across a period the ionization energies generally increase. Electrons in the same set of orbitals do not shield each other very well but the nuclear charge increases, making the electrons more difficult to remove. Z 3 4 Atom Li Be electron configuration 1s2 2s1 1s2 2s2 I1 (kJ/mol) 520 899 This trend is observed for Li and Be. The 1s electrons are the screening electrons for both Li and Be. The nuclear charge for Li is +3. The nuclear charge for Be is +4 and the 2s electrons in Be doesn't screen the nuclear charge very effectively. So an electron in Be is harder to remove than an electron in Li. However, there are some notable exceptions to this trend. Z Atom electron configuration I1 (kJ/mol) 2 1 3 Li 1s 2s 520 4 Be 1s2 2s2 899 2 2 1 5 B 1s 2s 2p 800 It is easier to remove an electron from B (Z = +5) than from Be (Z = +4). Why? (Hint: what type of electron is being removed?) For the next few elements the trend is followed. As Z increases, the I1 increases. Z Atom electron configuration I1 (kJ/mol) 3 Li 1s2 2s1 520 2 2 4 Be 1s 2s 899 5 B 1s22s2 2p1 800 2 2 2 6 C 1s 2s 2p 1090 7 N 1s2 2s2 2p3 1400 8 O 1s2 2s2 2p4 1310 There is another break between N and O. Question 1: (2 points) Fill in the energy diagram for each of these atoms: N __ __ __ __ __ 1s 2s 2px 2py 2pz O __ __ __ __ __ 1s 2s 2px 2py 2pz How do the electrons being removed differ from each other? The final two elements in this period follow the trend as Z increases, I1 increases. Z Atom electron configuration I1 (kJ/mol) 9 F 1s2 2s2 2p5 1680 2 2 6 10 Ne 1s 2s 2p 2080 More Questions: 2. Match the following electron configurations with the appropriate ionization energies (I1). (3 points) a) 1s2 2s2 3p6 3s2 3p6 4s2 3d10 4p6 5s1 b) 1s2 2s2 3p6 3s2 3p6 4s2 c) 1s2 2s2 3p6 3s2 3p6 4s2 3d104p6 i) 1356 kJ/mol ii) 595 kJ/mol iii) 409 kJ/mol 3. Which of the following would have the largest I1? Explain thoughtfully. (2 points) Na K Li Cs 4. What are the general periodic trends (i.e. trends as you go across a period) for: a. ionization energy? (1 pt) b. electronegativity? (1 pt) c. atomic radius? (1 pt) d. What is the major underlying concept that explains all of these trends? Explain carefully. (1 pts) 6. What is the relationship between: a. ionization energy and electronegativity? ( 1 pt) b. ionization energy and atomic radius (1 pt) c. electronegativity and atomic radius (1 pt) 7. Complete the table. (10 pts) Element Noble Gas Configuration Predicted charge on its ion Cs Cl Ga At S 8. Define Ionization Energy. (2 pts) _____________________________________________________________________________ _____________________________________________________________________________ 9. Write two separate equations to show the first and second ionization energy of calcium. (Remember state symbols are important as they from part of the definition). (2 pts) First Ionization _____________________________________________________ Second Ionization _____________________________________________________ 10. Which of the following elements (one from each pair) would you expect to have the highest first ionization energy? Explain your answers. Li and Fr (2 pts) _____________________________________________________ _____________________________________________________ K and Kr (2 pts) _____________________________________________________ _____________________________________________________ 11. Define the term, shielding. (2 pts) _______________________________________________________________________ _______________________________________________________________________ 12. Consider the table of the first four ionization energies for element A shown below. IONIZATION ENERGY IN KJ/MOL 1ST 578 2ND 1817 3RD 2745 4TH 11580 (i) In which group does A appear on the periodic table? _______________(1 pt) (ii) Predict the formula of the compound A forms with chlorine. ________ (1 pt) 13. (3 pts) Arrange the following species in order of increasing size. Rb+, Y3+, Br-, Kr, Sr2+ and Se2-. SMALLEST LARGEST ________________________________________________________________________ 14. Define isoelectronic? Give examples of two different cations and two different anions which are all isoelectronic. (3 pts) ________________________________________________________________________ ________________________________________________________________________ ________________________________________________________________________ 15. Define electron affinity. (2 pts) _______________________________________________________________________ ________________________________________________________________________ 16. Write an equation to summarize the process of second electron affinity of oxygen. (Remember state symbols are important as they from part of the definition). (2 pts) ________________________________________________________________________ 17. Consider the table of ionization energies for element X shown below. IONIZATION ENERGY IN 1ST 2ND 3RD 4TH 5TH KJ/MOL 496 4562 6912 9544 (i) In which group will X be found? (1) __________ (1 pt) (ii) Explain your answer. (2 pts) 13353 ________________________________________________________________________ ________________________________________________________________________ (iii) Predict the formula of X’s oxide. ___________ (1 pt) 18. Explain thoughtfully why magnesium tends only to form a +2 ion. (2 pts) _______________________________________________________________________ ________________________________________________________________________ 19. Why do elements in the same group react in similar ways? Explain well. (2 pts) ________________________________________________________________________ ________________________________________________________________________ 20. Identify any isoelectronic species in the following list. Fe2+, Sc3+, Ca2+, F-, Co2+, Co3+, Sr2+, Cu+, Zn2+ and Al3+. (3 pts) ________________________________________________________________________ 21. Arrange the following atoms into order of increasing first ionization energy. Sr, Cs, S, F and As. (3 pts) LOWEST HIGHEST ________________________________________________________________________ 22. For the following pairs of atom or ions, fill in the blank with either < or > to correctly order the identified periodic trend. (12 pts) Atomic radius Cd Ionic Radius Si __ Cl K+ __ Cs+ Br __ I F- __ Ne Au __ Hf Ar __ K+ Mg __ Al 3+ __ N3- Electron Affinity Cl __ I O __ F Na __ Mg 23. Fill in the blank with > or <. (3 pts) Atomic radius Ionization energy Na ______ S Na ______ S H __ He Electron affinity Na ______ S 24. What is a period? How many are there in the periodic table? (3 points) 25. What is a group (also called a family)? How many are there in the periodic table? (3 points) 26. State the number of valence electrons in an atom of: (4 points) a. oxygen b. barium c. iodine d. phosphorous 27. Name the following families: (3 points) GROUP FAMILY NAME GROUP1 GROUP 2 GROUP 17 GROUP 18 Noble gases 28. What is the name given to the group of elements that have the following valence shell electron configurations? (3 points) VALENCE SHELL s2 s2p6 s2p5 s1 ELECTRON CONFIGURATION Alkali FAMILY NAME metals 29. Which alkali metal belongs to the sixth period? (1 point) 30. Which halogen belongs to the fourth period? (1 point) 31. What element is in the fifth period and the eleventh group? (1 point) 32. For each of the following pairs, identify which of the two has the larger value. (6 points) CATEGORY ATOMS OR IONS WHICH HAS A LARGER VALUE? Electron affinity F or Cl Electron affinity S or Se Ionization Energy Na or Al Atomic Radius I or Br Ionic Radius Na+ or K+ Ionization Energy Fr or Ca