1. Semester I CP Chemistry Review 2009/2010

2. 1. What has to happen for a hypothesis to become a theory? It has to be supported by many, many experiments.

3. 2. Compare the size and mass of the nucleus in an atom. (where is most of the mass of an atom found?) The atom is mostly empty space. Although the nucleus occupies only a small part of the atom’s volume it contains most of its mass.

4. 3. All atoms of the same element have the same number of what subatomic particle? Protons

5. 4. What subatomic particles are in the nucleus and outside the nucleus? In nucleus = protons + neutrons Orbiting around nucleus = electrons

6. 5a. 131 I undergoes beta (β )emission 131 53 I  131 54 Xe + 0 -1 β (a loss of a beta particle = a gain of 1 in atomic # and a no change in mass)

7. 5b. 230 Th undergoes alpha(α) emission 230 90 Th  226 88 Ra + 4 2 α 4 2 α aka 4 2 He A loss of 4 atomic mass and a loss of 2 in atomic #

8. 5c. 230 90 Th undergoes gamma emission 230 90 Th  230 90 Th + gamma No change just release of energy

9. 6. Put the three types of decay (alpha, beta, gamma) in order from biggest to smallest. Then circle the type of decay that is the easiest to block. Alpha particles are the biggest and gamma the smallest Alpha easiest to block, gamma harder to block Alpha low energy and gamma high energy

10. 7. What are alpha particles? Beta particles? Gamma particles? Alpha are helium nuclei Beta are electrons Gamma are photons

11. 8. What occurs in a fusion reaction? What occurs in a fission reaction? Give an example of each. Fusion – two smaller nuclei combine to make one larger nucleus Fission – One large nucleus breaks down into 2 or + smaller nuclei

12. 9. How does the energy produced in a fission or fusion reactions compare to the energy produced in a chemical reaction? Both fission and fusion release A LOT more energy than a chemical reaction.

13. 18. What is the force that holds the nucleus together called? Strong nuclear forces 19. What is an isotope? Atoms with the same # of protons but different # of neutrons.

14. 20. How do you know if an isotope of an element is radioactive or stable? (remember there is ratio, one for elements #20 and under and one for those larger than #21) For elements #20 > the ratio of protons to neutrons is 1:1, then stable For elements #21 < the ratio of protons to neutrons is 1:1.5, then stable Or if in band of stability on graph (dots part)

15. 20.b Describe how an atom of Li would be changed for each of the following situations: a. Li atom loses or gains a neutron – loses or gains atomic mass b. Li atom loses or gains an electron – loses becomes a + ion, gains becomes a – ion. c. Li atom loses or gains a proton – becomes a different element

16. 22. Identify the element from the given electron configuration  1s2 2s2 2p6 3s2 3p5 – add p’s = 17 so Cl  1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 – p’s = 38, Sr  1s2 2s2 2p4 – p’s = 8, O

17. 23. Draw the electron dot (Lewis dot) diagram that indicates the number of valance electrons for #22 a, b and c. ClSr O

18. 24. Bromine (Br) commonly has a -1 charge while Potassium (K )has a +1 charge. Explain why this is and how it relates to the noble gas configuration. Why do ions commonly lose or gain ions to have a noble gas configuration? Br has 7 valence electrons and would like 8 to have a full outer orbital just like the noble gases, so it steals 1 electron away from another atom. K has 1 valence electron and would like to lose it so then it would be left with the next lower energy level which is full, just like a noble gas. So K likes to give its 1 valence electron away.

19. 25. Indicate which element has the greater atomic radius and give a reason for your choice. a. Na or Li- Na because atomic radius increases as you go down the group, plus Na has 3 energy levels and Li only has 2. b. Sr or Mg – Sr. Sr has 5 energy levels and Mg only has 3. c. C or Ge – Ge has 4 energy levels, C has 2 d. Se or O – Se has 4 energy levels, O has 2

20. 32. What is the difference between a group and a period on the periodic table? A group is a column and all in a group have the same # of valence electrons and behave similarly in a chemical reaction. A period is a row and tells you how many energy levels the element has.

21. 33. What is electronegativity and what is its trend down a group and across a row in the periodic table? The ability of an atom to attract electrons to itself. F has highest, Fr lowest Increases from left to right. Decreases as you go down a group.

22. 34. For each of the following pairs which atom/ion will have the HIGHER ionization energy? Explain. a. Ca or Ba? - Ca is higher because it has only 4 energy levels so its valence electrons are closer to its nucleus b. Al or Cl? – Cl, they both have 3 energy levels but Cl has more protons attracting its electrons

23. 35. Where are the metals, nonmetals and metalloids on the periodic table? Metals to the left, below staircase Nonmetals to the right, above staircase Metalloids along the staircase

24. 36. How many valence electrons does the atom have? a. Na – 1 (it’s in group IA) b. N – 5 (it’s in group VA) c. Ga - 3 (it’s in group IIIA) d. Mg – 12 (it’s in group IIA)

25. 37. What type of elements form ionic bonds ? Describe what happens to the electrons in an ionic bond? Why do the ions in an ionic compound stay together? Ionic – metal with nonmetal Electrons are transferred from metal to nonmetal Metal becomes + and nonmetal becomes –. Opposites attract. Electrostatic forces.

26. 36. What type of elements form covalent bonds? What are the electrons doing in a covalent bond? 2 or more nonmetals Electrons are being shared between the atoms

27. 37. Describe what the electrons are doing in a metallic bond. What types of atoms form metallic bonds? Sea of electrons. Valence electrons are floating from one atom to another. Metallic bonds form only between metal atoms!!

28. 38. What type of ions to metals form? How about non-metals? Metals – cations (+ ions) Nonmetal – anions (- ions) 39. What is a coefficient? The number before a formula in a chemical equation. It tells you how many atoms or molecules of that substance is involved in the reaction.

29. 40. Balance the following equations  2 Al + 2 H3PO4  2AlPO4 + 3 H2  __P4 + __S8  __ P4S8  2 C2H6 + 7 O2  4 CO2 + 6 H2O  2 AgNO3 + ___K2S  ___Ag2S + 2 KNO3  ___H2CO3  ____H2O + ____CO2

30. 40b. Why must chemical equations be balanced? Because matter cannot be created nor destroyed in a chemical reaction. So the numbers of elements in the reactants must equal the numbers of elements in the products

31. 41. What is the mass in grams of each of the following: a. 1 mole of Al= 27 g b. 0.5 moles of S = 16 g c. 1 mole of NaOH= 40 g (Na = 23 + o = 16 + H = 1) d. 3 moles of O2 = 32 g (O = 16 x 2) You must use the atomic masses from the periodic table

32. 42. How many moles are in the following: b. 22 grams of N2O 22g N2O 1 mole N2O = 0.5 mole 44g N2O b. 342 grams of Al2(SO4)3 342 g Al2(SO4)3 1 mole Al2(SO4)3 = 1 mole 342 g Al2(SO4)3 c. 10 grams of Fe2O3 10g Fe2O31 mole Fe2O3 = 0.06 mole 160 g Fe2O3

33. 43. How many atoms of molecules are in the following? a. 2 moles of sugar – 1.2 x 10 24 b. 0.5 moles of H2O – 3.0 x 10 23 c. 42 grams of C3H6 42g C3H6 1 mole C3H6 6.0 x 10 24 molecules 42 g C3H61 mole C3H6 = 6.02 x 10 23

34. 43b. How many atoms in a. 32 g of O 32g O 1mole O 6.02 x 10 23 atoms 16 g O 1 mole O = 1.2 x 10 24 atoms b. 23 g of Na 23g Na 1mole Na 6.02 x 10 23 atoms 23g Na1 moleNa = 6.02 x 10 23 atoms

35. c. 36 g of C 36g C 1mole C 6.02 x 10 23 atoms 12 g C 1 mole C = 1.8 x 10 24 atoms

36. 43c. How many atoms in 1 mole of Br? = 6.02 x 10 23 atoms How many molecules in 1 mole of Fe2O3? = 6.02 x 10 23 molecules

37. 44. Calculate the volume of these gases at STP a. 1.5 moles of Neon 1.5 mole Ne 22. 4 L = 33.6 L 1 mole Ne b. 32 grams of oxygen 32g O 1 mole O 22.4 L = 44.8 L 16g O 1 mole O

38. 48. Given the following balanced equation: 2NH3 + 3O2 + 2CH4  2HCN + 6H2O A. How many moles of HCN are produced from 9 moles of O2? 9 mole O2 2 mole HCN = 6 moles HCN 3 mole O2

39. B. How many moles of NH3 are needed to produce 24 moles of H2O? 24 moles H2O 2 moles NH3 = 8 moles H2O 4 moles of H2O

40. 49. Given the following balanced equation: SiO2 + 4HF  SiF4 + 2H2O a. If 5 grams of HF reacts with SiO2, what mass of H2O forms? 5g HF 1 mole HF 2mole H2O 18g H2O 20 g HF4 mole HF 1 mole H2O = 2.25 g H2O

41. b. If 120 grams of SiO2 reacts with excess HF, how many grams of water will form? 120g SiO2 1mole SiO2 2mole H2O 18g H2O 60g SiO2 1 mole SiO2 1mole H2O = 72 g H2O