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# Final Review

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### Final Review

1. 1. Significant Figures a.k.a.- sig figs
2. 2. Significant Digits <ul><li>The certain digits and one estimated digit of each measurement are significant. </li></ul><ul><li>Remember! Every time you make a measurement, you record all of the certain digits and one estimated digit. </li></ul>200.5 4 g
3. 3. Rules for Sig Figs <ul><li>Non zeros are always significant. </li></ul><ul><li>Zeros between non zeros are significant. </li></ul><ul><li>Zeros at the end of significant digits following a decimal point are significant. </li></ul><ul><li>*They show precision in measurement. </li></ul><ul><li>4) Place keeper zeros are NOT significant. </li></ul><ul><ul><li>Zeros preceding significant digits. </li></ul></ul><ul><ul><li>Zeros following significant digits without a decimal point. </li></ul></ul>
4. 4. Try These Examples <ul><li>7.05940 </li></ul><ul><li>Final zero significant (follows decimal point) </li></ul><ul><li>6 significant digits </li></ul><ul><li>0.00135 </li></ul><ul><li>Leading zeros Not significant (place keepers) </li></ul><ul><li>3 significant digits </li></ul><ul><li>20,400 </li></ul><ul><li>Final zeros Not significant </li></ul><ul><li>(place keepers – no decimal) </li></ul><ul><li>3 significant digits </li></ul>
5. 5. Heat Calculations <ul><li>Heat (q) </li></ul><ul><ul><li>Energy transferred from an object at a higher temperature to an object at a lower temperature. (heat lost = -heat gained) </li></ul></ul><ul><ul><li>q = mc  T </li></ul></ul><ul><ul><li>q=mH fus </li></ul></ul><ul><ul><li>q=mH vap </li></ul></ul>
6. 6. Heat Calculations <ul><li>A 10.0g sample of iron at 50.4 o C is cooled to 25.0 o C in 50.0g of water. Calculate the amount of heat lost by the iron. </li></ul><ul><li>c iron= 0.449 J/g o C </li></ul><ul><li>A 2.1g ice cube at –8.0 o C melts completely and warms to 12.5 o C. How much heat was required? </li></ul><ul><li>H fus ice = 334 J/g </li></ul><ul><li>c ice = 2.03 J/g o C </li></ul><ul><li>c water = 4.18J/g o C </li></ul>
7. 7. <ul><li>soluble – a substance that dissolves in a solvent </li></ul><ul><li>insoluble – a substance that does not dissolve in a solvent </li></ul><ul><li>solvation – the process of surrounding solute particles with solvent particles to form a solution </li></ul>
8. 8. Solvation <ul><li>When a solid solute is placed in a solvent, the solvent particles completely surround the surface of the solid solute. </li></ul><ul><li>If attractive forces between the solute particles and the solvent are greater than the attractive forces holding the the solute particles together, the solvent particles pull the solute particles apart and surround them. </li></ul>
9. 9. + - - - NaCl Na = Cl = H 2 O H = O = + - + Process of Solvation - + + - + + + - - + + - + - - + - + - + - + - + + - + + - + + - + + - + +
10. 10. Water- Universal Solvent <ul><li>Polar molecule </li></ul><ul><li>Dipoles allow solvation of ions and polar molecules </li></ul>
11. 11. Factors that Affect the Rate of Solvation <ul><li>Agitate the solution (stirring) </li></ul><ul><li>increase the temperature of solvent </li></ul><ul><li>increase the surface area of the solute </li></ul>
12. 12. SOLUBILITY – refers to the maximum amount of solute that will dissolve in a given amount of solvent (at a specified temperature and pressure) <ul><li>saturated – a solution that contains the maximum amount of dissolved solute (for a given temp & pressure than a saturated solution) </li></ul><ul><li>unsaturated – a solution that contains less dissolved solute (for a given temp & pressure than a saturated solution) </li></ul>
13. 13. Units of Solubility <ul><li>g of solute </li></ul><ul><li> 100 g water </li></ul><ul><li>read these units as: </li></ul><ul><li>“ grams of solute per 100 grams of water ” </li></ul>
14. 15. Here are some for you to try. <ul><li>What mass of solute will dissolve in 100mL of water at the following temperatures. Also determine which of the three substances is most soluble in water at 15°C. </li></ul><ul><li>KNO3at 70°C </li></ul><ul><li>NaCl at 100°C </li></ul><ul><li>NH4Cl at 90°C   </li></ul>
15. 16. <ul><li>supersaturated – a solution that contains more dissolved solute than a saturated solution at the same temperature </li></ul><ul><li>- is above the solubility curve </li></ul>solute will usually precipitate out of solution
16. 17. Radioactivity <ul><li>There are two main types of radioactivity: Natural and Induced </li></ul>
17. 18. Natural Radioactivity <ul><li>Occurs in nature </li></ul><ul><li>Usually large, unstable nuclei </li></ul><ul><li>Occurs in three ways: </li></ul><ul><ul><li> Particle (alpha particle) </li></ul></ul><ul><ul><li> Particle (beta particle) </li></ul></ul><ul><ul><li> Ray (gamma ray) </li></ul></ul>
18. 19. Alpha Decay <ul><li>A helium nucleus is released from the nucleus. ( ) </li></ul><ul><ul><li>The mass decreases by 4 </li></ul></ul><ul><ul><li>The atomic number decreases by 2 </li></ul></ul><ul><ul><li>(Because the He nucleus has 2p + and 2n o ) </li></ul></ul><ul><li>Alpha radiation can be stopped by a piece of paper. Cannot penetrate skin. Not dangerous. </li></ul>
19. 20. Alpha Decay Example Notice that the uranium has changed into a new element, thorium.
20. 21. Beta Decay <ul><li>An electron is released from the nucleus when a neutron becomes a proton. </li></ul><ul><li>The mass is unaffected. (the mass of a neutron is roughly equal to the mass of a proton) </li></ul><ul><li>The atomic number is increased by 1. </li></ul><ul><li>Harder to stop and more dangerous. </li></ul>
21. 22. Beta Decay Example Notice that carbon has changed into nitrogen.
22. 23. Gamma Decay <ul><li>Pure energy is released from the nucleus. </li></ul><ul><li>The mass and atomic number are unaffected. </li></ul><ul><li>Stopped by lead. The most harmful to living tissue. </li></ul>
23. 24. Gamma Decay Example No new element formed. Gamma radiation (energy) released.
24. 25. Induced Radioactivity <ul><li>Particles are slammed together to cause transmutation of stable elements. (Nuclear Bombardment) </li></ul><ul><li>Discovered by Rutherford in 1919. </li></ul>
25. 26. Uranium-238 Decay Series
26. 27. Average Atomic Mass (How the number ends up on the periodic table!!) 1 st  Mass of one isotope x % abundance in decimal form (watch SIG FIGS!!) 2 nd  Do this for each isotope of that element 3 rd  Add all individual isotopes together to get the average atomic mass of the element.
27. 28. 1. Calculate the average atomic mass of potassium using the following data: Potassium-39 38.964 amu x 0.9312 = 36.28 amu Potassium-41 40.962 amu x 0.0688 2.82 amu = + Average atomic mass for K = 39.10 amu 6.88 % 40.962 amu Potassium-41 93.12% 38.964 amu Potassium-39 % abundance Mass Isotope
28. 29. 2. Calculate the average atomic mass of magnesium using the following data: Magnesium-24 23.985 amu x 0.7870 = 18.88 amu Magnesium-25 24.986 amu x 0.1013 2.531 amu = + Average atomic mass for K = 24.31 amu + Magnesium-26 25.983 amu x 0.1117 = 2.902 amu 11.17 % 25.983 amu Magnesium-26 10.13 % 24.986 amu Magnesium-25 78.70% 23.985 amu Magnesium-24 % abundance Mass Isotope