PHYSICAL SCIENCE SLT STUDY GUIDE Chemistry and Physics 2012-2013

1. PHYSICAL SCIENCE SLT STUDY GUIDE Chemistry and Physics 2012-2013

2. How to convert between American and SI • Dimensional Analysis (also called Factor-Label Method or the Unit Factor Method) is a problem-solving method that uses the fact that any number or expression can be multiplied by one without changing its value. It is a useful technique. • Unit factors may be made from any two terms that describe the same or equivalent "amounts" of what we are interested in. For example, we know that 1 inch = 2.54 centimeters • We can make two unit factors, each equaling 1 over 1 (or 1) from this information:

3. Now, we can solve some problems. Set up each problem by writing down what you need to find with a question mark. Then set it equal to the information that you are given. The problem is solved by multiplying the given data and its units by the appropriate unit factors so that only the desired units are present at the end. • For example: You want meters, and are given the data in inches. So – Given Unit (inches) x Wanted Unit (meters) Equal Amount in Given Units (# inches in a meter) The inches cancel-out leaving your answer in meters

4. Example: • Note the fraction you multiply by is equal to one over one. (1 inch = 2.54 cm) = You use an one/one fraction each time regardless of what the conversion is! = This is why learning the basic equivalences will make life easier

5. More examples: How many seconds in two years? • What is the density (D) of mercury (13.6 g/cm3) in units of kg/m3?

6. ATOMIC MODELS Protons (+ charge) and Neutrons (0 charge) make up the nucleus; Electrons (- charge) around nucleus and held by electromagnetic force. p+ and n made up of quarks. Plum Pudding Model Electron Cloud Model

7. Four Universal Forces

8. ISOTOPES • All atoms of an element have the SAME number of protons (p+) • The p+ number is the atomic number (Z) • This is a constant – it stays the same for that element’s atoms • For example: All Sodium (Na) atoms have 11 p+ • If an atom loses a proton, it becomes a different element • If Na loses 1 p+, then it has become Neon (Ne)

9. Z = atomic number = p+ • The number of protons identifies the atom and which element it is • In a stable atom: • # p+ = # n0 = # e- • Thus, Na in its stable form has 11 p+; 11 n0; and 11 e- • If it has an unequal number of p+ and n0, then it is called an ISOTOPE

10. The Carbon Isotope

11. IONS • Ions are when an atom has an unequal number of p+ and e- • Metals form (+) ions and nonmetals (-) ions • Remember – a stable atom has a neutral overall charge due its equal number of p+ and e- • When an atom loses or gains an e-, its charge changes accordingly • Loss of e- means a + charge; gaining an e- means a – charge for the atom

12. Losing or Gaining e- . . . . . • If an atom loses an e-, then it has more p+ than e- and it will have an overall positive charge • Different elements’ atoms can lose 1, 2, 3, or even 4 electrons depending on various factors • If an atom has LOST e-, then it is called a CATION or a positive ion • A Cation would be written as Al+ (the one being understood) or Al+3

13. Atoms can also gain electrons • If an atom gains electrons (from 1 up to 4), then it will have more e- than p+ and will end up having an overall negative charge • A negatively charged ion is called an ANION • (A positively charged ion is called a CATION) • The NOBLE GASES will not form ions and thus will not bond • The Transition Metals can form various numbers of positive ions – got to learn these! • The losing or gaining of electrons determines what type of bonds the atoms will form, and which atoms will bond to others

14. Using the Periodic Table • Elements in the Main Groups (A), form fairly consistent ions • Group IA will form +1 ions; Group 2A form up to +2; Group 3A form up to +3 ions • Group 4A will form either up to -4 or +4 ions • Group 5A will form up to -3 ions; Group 6A up to -2; Group 7A form -1; and Group 8A will not form ions at all • Those elements in the B group (transition metals) vary in their + charges meaning they can form different ions

15. Group Names - Periodic Table

16. Know These!

17. Ions and Isotopes in Review • Stable atom: #p+ = #n0 = #e- • Atomic Mass: #n0 = # p+ • If charge is 0, then #p+ = #e- • If charge is positive, then #p+ > #e- Cation • If charge is negative, then #p+ < #e- Anion • Mass measured in AMU (Atomic Mass Units) based on the C-12 atom

18. Examples: • Li-1 has gained an electron, meaning there is one more negative charge than positive ones • It has 3 p+ and 4 e- • Li+1 has lost an electron, meaning there is one more positive charge than negative ones • It has 3 p+ and 2 e- • REMEMBER: The # of p+ DOES NOT CHANGE • Only the number of n0 (isotope) and e- (ion) change

19. Cf-3 has an atomic number of 98 • This means it has 98 p+ • Its atomic mass is 216 • It has 118 n0, (216 – 98), making it an ion and an isotope! • Since it has a -3 charge, the number of e- will be 101; (98 + 3) • Zn+1 has 30 p+ and n0; but due to the +1 charge, it has only 29 e-

20. On the Periodic Table: The top number is Z, the Atomic Number or number of p+ The Element’s Symbol The element average atomic weight set by isotopes and abundances

21. Counting Atoms in a Molecule In the example, NH3, the subscript 3 only applies to the hydrogen. • Therefore: there is 1 N and 3 H in ammonia In the example, 3Ca3(PO4)2, the number of atoms changes due to the Coefficientis always in front of the whole molecule!! -The subscript 2, multiplies the P (2) and O (4 x 2 = 8) since it is outside the parenthesis -The subscript 3 only goes with the Ca -The coefficient 3 is multiplied to the Ca, P and O after you do the subscripts -Therefore, this molecule has (3 x 3) Ca+ (3 x 2) P + (3 x 4 x 2) 0 which equals 39 atoms 3Ca3(PO4)2

22. Ionic Bonds • These are the bonds between a metal and a nonmetal • The metal Ion is positively charged and called a cation • The nonmetal Ion is negatively charged and called an anion • The bonded molecule should be neutrally charged when finished

23. Knowing where the metals and nonmetals are on the table will make your life easier

24. Covalent Compounds • These can be monatomic or polyatomic compounds • It is a bond between two nonmetals • They share a pair of electrons • They can be subgrouped into polar or nonpolar • If a binary compound (2 atoms) – use the same naming rules as in Ionic Compounds

25. Naming Covalent Compounds Process: • Prefix Indicating # + full name of first nonmetal • Prefix Indicating # + root name of second nonmetal + suffix “ide” • Watch for polyatomics and use their proper names

26. If it has more than two atoms – need to use the prefixes Number PrefixNumber Prefix 1 Mono 7Hepta 2 Di 8Octa 3 Tri 9 Nona 4 Tetra 10Deca 5 Penta11Undeca 6 Hexa12Dodeca

27. For Example: NOTE THESE ARE ALL NONMETALS WITH NONMETALS! • P4S10becomes TetraphosphorousDecasulfide • P2O5becomes DiphosphorousPentaoxide • SF6 becomes Sulfur Hexafluoride • N2O3 is Dinitrogen Trioxide • CO is Carbon Monoxide • SO2 is Sulfur Dioxide • SiBr4 becomes Silicon Tetrabromide • Water is really Dihydrogen Monoxide!

28. Naming Ionic Compounds is really simple: 1. Name the cation (metal) using its proper name; if it is a polyatomic, do the same 2. Then, using the stem of the anion (nonmetal), simply add the suffix “ide” to it 3. If a transition metal with different possible ions, a roman numeral will tell you which one it is – and it changes the molecular formula! Examples: Iron (II) Sulfide = Fe+2 and S-2 combined Zinc + Chlorine = Zinc Chloride Iron + Oxygen = Iron Oxide Lithium + Cyanide = Lithium Cyanide Ammonium + Fluorine = Ammonium Fluoride Cobalt + Phosphorous = Cobalt Phosphide

29. Balancing Compounds In an Ionic Compound – balance the molecule using the criss-crossrule. Switch oxidation numbers, making them into subscripts and DROP charges. Mg +2 + Cl-1 Mg Cl2 The one is understood. This applies even if using a polyatomic ion

30. NH4+ + O-2 (NH4)2O The parentheses are used to keep the polyatomic together; the 1 is understood Pb+4+ CO3-2 Pb2 (CO3)4 and this can be simplified by reducing the subscripts to Pb(CO3)2

31. Chemical Equations • The chemical equation is the rxn formula • Reactants  Products • Each component will have a phase indicator: • (g) meaning it is in its gaseous phase (not just gassy) • (l) meaning it is in its liquid phase • (s) in its solid phase • And (aq) meaning the substance is in a solution of water, aq meaning aqueous

32. Must remember which elements are normally diatomic (N2, O2, F2, Cl2, Br2, I2, and H2) • All molecules in an equation must be balanced first!! • Remember the criss-cross rule!! • You may not adjust any subscripts from the original formula • You may add and adjust, as you will see, the coefficients in front of each item in the equation

33. Example: • H2 + O2 H2O • This is the skeleton equation • According to the Law of Conservation of Matter, both sides of the arrow must have the SAME number of atoms for each and every element – NO EXCEPTIONS • The  can be treated like an = sign • In reality, it indicates that some sort of process occurred to cause the reaction • So. . . . .

34. To balance this simple equation: • We ARE NOT ALLOWED TO CHANGE SUBSCRIPTS • We CAN ADJUST COEFFICIENTS ONLY • The subscripts are the numbers after and below each element’s symbol • The coefficients are number in front of a unit (atoms or molecules) and tell how many units there are • The coefficients are multiplied out to each and every unit’s atom they are in front of • So. . . . .

35. H2 + O2 H2O • There are 2 H and 2 O on the reactant side of the equation (the left side) • There are 2 H and only 1 O on the product side (the right side) • Each side must balance • You may add, adjust, finagle, cram, etc. any coefficient in front of any and/or all units to get the equation to balance • Therefore: 2H2 + O2 2 H2O

36. 2H2 + O2 2H2O • Now this is balanced! • It means it takes 2 hydrogen molecules and one oxygen molecule to form 2 water molecules

37. Balancing Equations Steps: • First identify all the reactants and products in the equations • Remember – subscripts indicate how many of each element’s atoms are present – with 1 being understood • Remember to multiply out all subscripts that are outside a unit in parentheses! • YOU CAN’T CHANGE SUBSCRIPTS • COEFFICIENTS HAVE TO GO IN FRONT OF A UNIT

38. Let’s take the unbalanced equation of: KClO3 KCl + O2 • List the elements and how many for both sides of the arrow K 1  K 1 Cl 1 Cl 1 O 3 O 2 • Obviously, everything is fine except for oxygen • This is where we have to adjust

39. We can only use coefficients • So we try to multiply each Oxygen by a number to get them to equal out • These multipliers become coefficients K 1  K 1 Cl 1 Cl 1 O 3 x 2 = 6 O 2 x 3 = 6 • So the new equation is: 2 KClO3  KCl + 3 O2 • This changes the number of K and Cl now • You have to readjust again. . . . . .

40. 2 KClO3 KCl + 3 O2 • Now we have: K 2  K 1 Cl 2 Cl 1 O 6 O 6 • Multiply the product KCl by a coefficient of 2 and it balances • Let’s check: 2 KClO3  2KCl + 3O2 K 2  K 2 Cl 2 Cl 2 O 6 O 6 • It’s Balanced! Finally.

41. Another Example: C2H6 + O2 CO2 + H2O • List the atoms and numbers: C 2 C 1 H 6 H 2 O 2 O 2 + 1 = 3 • Let’s go with C first by multiplying CO2 by a coefficient of 2 C2H6 + O2  2 CO2 + H2O

42. This gives us: C 2 C 1 x 2 = 2 H 6 H 2 O 2 O 4 + 1 = 5 • Now, let’s balance H by multiplying H2O by 3 • This gives us: C2H6 + O2 2 CO2 + 3 H2O C 2 C 2 H 6 H 2 x 3 = 6 O 2 O 4 + 3 = 7 • It’s still not balanced! • Let’s try readjusting Oxygen to get it the same amount

43. So, if we change the reactant oxygen to 7 and the product water to 6, we get: C2H6 + 7 O2 2 CO2 + 6 H2O • This also changes our product hydrogen. • Therefore, change the reactant C2H6 and the product CO2 to balance and you get: 2 C2H6 + 7O2  4 CO2 + 6 H2O C 4 C 4 H 12 H 12 O 14 O 14

44. Reaction Types • SYNTHESIS (or Direct Combination or Composition) REACTIONS • 2 + reactants join together to form a single product • Resulting compound is based on common oxidation numbers of the reactant elements • There is typically an electron transfer from the element with the lower EN to the one with the higher EN • So: A + B  AB or AB + C  ABC

45. If two nonmetals involved – a covalent bond formed • If two metals – a metallic bond • If metal with a nonmetal – ionic bond

46. DECOMPOSITION REACTIONS • Compounds break down into components • AB  A + B or ABC  AB + C • Examples. . . CaCO3 CaO + CO2 Ca(OH)2 CaO + H2O 2 KClO3 2 KCl + 3 O2 2 NaCl 2 Na + Cl2 H2CO3 H2O + CO2

47. REPLACEMENT REACTIONS (2 types) • Single Replacement (Displacement Rxn) • Key Rule: Metals Replace Metals • A + BC  AC + B • If Nonmetal – a transfer of e- from more reactive to lesser one • Halogens Replace Halogens also • Metals replace H in H2O  Metal OH- + H2 (g) • Metals replace H in Acids  salt + H2(g) • Al + H2SO4  AlSO4 + H2(g) • 2 Sc(s) + 6 HCl (aq)  2 ScCl2(aq) + 3 H2(g)

48. DOUBLE REPLACEMENT Example: • FeCl3 + 3 NaOH 3 NaCl + Fe(OH)3 OH goes with FE Cl goes with Na • Cations exchange anions with each other • No change in oxidation numbers • Better know your ions and polyatomics • Remember the criss-cross rule and balance each compound after exchanging anions! • So: AB + CD  AC + BD

49. COMBUSTION • An exothermic rxn (gives off energy) • Usually find CO2 and H2O in products • O2 usually found in reactants • CH4(g) + 2 O2 CO2(g) + 2 H2O(g) + heat • 2 C4H10(g) + 13 O2(g)  8 CO2(g) + 10 H2O(g)

50. ACID/BASE REACTIONS An acid + base  salt + H2O • Acids lose a H+ ion and the bases lose OH- ion • These make up one of the products, water • Process is called neutralization • The produced salt does not have to be NaCl and can be any ionic compound • Measure acid with pH scale (1 strong, 7 neutral and 14 is a base) • Measure base with pOH scale