Preparation for STAR

1. Preparation for STAR Important things to remember

2. Positively charged particle Negatively charged particle Neutral particle Positive dense core #neutrons + #protons # protons in an atom Proton Electron Neutron Nucleus Mass number Atomic number Structure of the atom

3. Positively charged ion Negatively charged ion Charge on any atom Discovery of electron Discovery of nucleus Cation Anion None JJ Thomson Rutherford

4. 6.2

5. 6.2 Transition Elements • Blocks of Elements

6. 6.3 Trends in Atomic Size

7. 6.3 Trends in Atomic Size

8. 6.3 Trends in Ionization Energy

9. 6.3 Trends in Ionization Energy

10. 6.3 Trends in Ionic Size • Relative Sizes (in pm) of some atoms and ions

11. 6.3 Trends in Ionic Size • Trends in Ionic Size Size generally increases

12. 6.3 Trends in Electronegativity

13. Nomenclature • Memorize common acids pg 105 • Practice naming type I, II, III cpds.

14. Na2SO4 CO2 Fe(NO3)2 copper(II) chloride boron trifluoride H2SO4 HNO3 NH4NO3 acetic acid sodium sulfate carbon dioxide iron(II) nitrate CuCl2 BF3 sulfuric acid nitric acid ammonium nitrate CH3COOH Practice

15. Types of reactions • Mg + Cl2  MgCl2 • synthesis or combination • 2H2O  2H2 + O2 • decomposition • BaCl2 + K2CO3 BaCO3(s) + 2KCl • double displacement • Zn + 2HCl  ZnCl2 + H2 • single replacement

16. Matter • anything that occupies space and has mass • elements cannot be broken into simpler substances • Solids have a lot of intermolecular forces • Liquids have intermolecular forces and also the ability to flow • Gases have very little intermolecular forces

17. three ways to make dissolving faster another name for a homogeneous mixture substance present in large quantities in a mixture substance present in small quantities moles /liter stirring, heating surface area solution solvent solute molarity Solubility and Solutions

18. bond formed by equal sharing electrons bond formed by unequal sharing electrons bond formed by transferring electrons non-polar covalent O2 polar covalent HF ionic MgCl2 Bonds

19. Lewis structure for carbon atom CO2 Lewis structures .. .. :O=C=O:

20. Moles • The atomic mass of all elements on the periodic table is the mass of 1 mole of that element where one mole equals_________ • 6.022 x 1023atoms • How many atoms of barium are there in 2.09g of barium? • 9.16 x 1021 atoms of Ba • How many atoms of oxygen is present in 32.00g of oxygen? • 1.2 x 1024 atoms of oxygen

21. What is the mass% of each element in isopropyl alcohol C3H7OH • 13.42 %H, 59.96% C,26.63% O • When iron(III) oxide reacts with aluminum metal it produces 0.625moles of aluminum oxide. If iron is the other product how much of iron was produced? Fe2O3 + Al- Al2O3 + Fe • 1.25mol Fe

22. In the reaction, TiCl4 + O2 TiO2 + Cl2 theoretical yield of TiO2 is 2.83g. If actual yield is 2.12g, what is the percent yield? • 75% • What is the empirical formula for the compound C6H6? • CH • How many grams of N2 are formed when 18.1g of NH3 reacts with 90.4g of CuO? • 2NH3 + 3CuO  N2 + 3Cu + 3H2O • 10.6gN2

23. H2O (g) H2O (l) + energy energy + 2HgO (s) 2Hg (l) + O2 (g) energy + H2O (s) H2O (l) Exothermic process is any process that gives off heat – transfers thermal energy from the system to the surroundings. Endothermic process is any process in which heat has to be supplied to the system from the surroundings. 6.2

24. Energy Diagrams Exothermic Reaction 100kJ produced

25. Endothermic Reaction 200kJ absorbed

26. The heat capacity (C) of a substance is the amount of heat (q) required to raise the temperature of a given quantity (m) of the substance by one degree Celsius. Heat (q) absorbed or released: q = msDt s = specific heat capacity of that substance Dt = tfinal - tinitial m =mass in grams 6.5

27. 11.8

28. Convert 25ºC to Kelvin STP conditions What is molar volume? At const. T when pressure increases, volume_____ At const. P when temperature increases, volume __________ absolute zero the higher the altitude lower the _________ 298 K 1atm, 273K 22.4L decreases increases 0K or -273ºC atmospheric pressure Gas laws

29. Equilibrium • Chemical equilibrium is achieved when: • the rates of the forward and reverse reactions are equal and the concentrations of the reactants and products remain constant • Le Chateliar’s principle states that when you have a system in equilibrium, the following rules apply: • Adding a substance will cause the equilibrium to shift away from the substance • Removing a substance will cause the equilibrium to shift toward the substance • Increased pressure will cause the eqm. to shift to the side with smaller number of gas molecules. • decreased pressure will cause the eqm. to shift to the side with greater number of gas molecules.

30. Alpha, Beta, and Gamma • Historically, the products of radioactivity were called alpha, beta, and gamma when it was found that they could be analyzed into three distinct species by either a magnetic field or an electric field.

31. 222 218 88 86 4 4 2 2 Types of Radioactive decay • Alpha particle emission α-particle is a helium nucleus ( He) and this type of emission is very common mode of decay for heavy radioactive nuclides. By releasing the alpha particle, there is an increase in n/p ratio. Isotopes with Z>83 undergo alpha emission. example: Ra  He + Rn

32. 0 1 0 1 0 3 -1 0 -1 1 -1 1 3 2 Beta particle emission (β-particle)e • This is another common form of decay that results in the production of a very light particle with Z= -1 and mass# A = 0. This happens in light nuclei with too many neutrons (n/p decreases). Example: H  e + He note: the net effect of β-particle emission is the conversion of a neutron to a proton. n  e + H