1. PHYSICAL SCIENCE SLT STUDY GUIDE Chemistry and Physics 2014-2015

2. Mass vs. Matter Matter is something that occupies spacetime and has length, width and height (and time) Mass, on the other hand, is HOW MUCH MATTER an object has at any one time – It is a measure of inertia – Measured in grams, milligrams or kilograms, etc. – The mass of an object does not change – Mass and Matter follow the Conservation Laws – they can’t be created nor destroyed – only transformed

3. Mass remains the same regardless of the phase of the matter However – the volume may change due to the substances density in that state of matter 1 kg  1 kg  1 kg Small VolumeMedium Volume High Volume High DensityMedium Density Low Density

4. NOTE: Mass is NOT the same as Weight Weight is mass with the effect of gravity applied to it! – Wt. = mass x gravity (9.81 m/s 2 on Earth) Mass does not change – gravity does change from place to place – so weight will change

5. Volume and Density Volume is how much space a certain mass occupies – typically calculated by multiplying width x height x length of the container – Measured in cm 3, L, ml, etc. Density is how much mass is compacted into a space – In other words, how closely packed are the atoms of the matter – D = mass / volume – Answers will be in g/ml, g/L, etc.

6. Types/Phases of Matter The type or phase of matter is dependent upon the KE in the particle’s of that matter All particles (atoms or molecules) have KE They move by rotating, vibrating or translating (going in a straight line) So, the amount of KE determines the phase of the matter – Particle KE is affected by things like temperature and pressure

8. 1. PLASMA On the scale of KE, this phase/type of matter has the highest amount In fact, it has so much, the electrons get ripped off the atom’s nucleus – leaving plasma Example is lightning or the surface of the sun

9. 2. Gas Gases are next on the KE scale A gas is defined as a type of matter that fills its container, so its volume equals whatever it’s in Gases expand to fill their container in a process called diffusion Since the atoms of a gas have a lot of space between them, gases can be condensed or compressed – In fact – applying pressure and lowering the temp of a gas will transform it into a liquid

10. Gas particles – due to the space and KE – collide a lot – They collide with each other and the sides of the container they are in – The collisions with the container we measure and call pressure – The rate of collision is the gas pressure in the container!

11. 3. Liquid Like gases, they fill the container they are in – so their volume is measured by their container – A liquid’s mass and volume can be measured separately However – there is less space between the particles and they are less compressible

12. Key Terms with Liquids Viscosity – this is the rate or ability of a liquid to flow Surface Tension – the top of a liquid has a “skin” on top due to the attraction of the top layer of particles by the ones below it – This is what allows a drop to form a circular shape

13. Vapor Pressure - due to evaporation, there will be some particles of a liquid in the atmosphere above it, and this changes the pressure above that liquid Capillary Action – this is the ability of a liquid to flow uphill! – It is achieved through electromagnetic forces!

14. 4. Solids Much lower particle KE and this is limited mainly to vibrations and rotations Solids have a set volume, mass, shape and density Just like the other types of matter – solids are affected by temperature and pressure

15. 5. Bose-Einstein Condensate This form of matter can only be achieved at absolute zero! – That’s -465 o F, -273 o C, or O o K!!!!! This means that there is NO KE in or between the matter’s particles Because of this – light sent through BEC will slow down to 7, that’s right, 7 mph!!!!! Temp lowered by using lasers

16. Phase Diagrams and Concepts Phase diagram – a graph to show the phases of a substance and at what temperature and pressure it goes into each phase

17. Phase Changes: From  To Term Solid LiquidLiquification / Melting LiquidSolidSolidification / Freezing LiquidGas/VaporEvaporation Gas/VaporLiquidCondensation SolidGasSublimation GasSolidDeposition / Crystallization

18. Classifying Matter Baryonic Matter (Observable Matter – 20% Universe) SubstancesMixtures ElementCompound Heterogeneous Homogeneous -On Periodic-2+ elements -Separable by -Includes: Table-Molecules filtering, distillingColloids and -Pure Form-Proportional or crystalizingSolutions -Single group make up -Suspension-Non-separable of Atoms -Emulsion

19. Matter Concepts: Element – The pure form of a substance – Found on the periodic table – Can be in any of the phases mentioned! Compound – A combination of 2 or more elements – Combination is in a ratio or proportional – Basically – a molecule

20. Mixtures – 2 or more elements or compounds mixed together – Heterogeneous – a mixture that can be easily separated through filtration, distillation or crystallization – Suspension – a temporary mixture where the particles separate out on their own (like dirt in water) – Emulsion – a suspension of two liquids – will separate (mayonnaise)

21. – Homogeneous Mixture – non-separable, remain mixed completely and equally – Includes Colloids – a solution with large particles dissolved in it, like milk or paint – Includes Solutions – a solution is made up of the solute (what is dissolved) and the solvent (what the solute is dissolved in) Example: Salt Water  Salt is the Solute and water is the solvent

22. Chemical and Physical Properties PhysicalChemical Does not change with phaseMay change in a reaction Density, MassAtomic Structure Boiling/Freezing PointsFlammability ColorOxidation HardnessReactivity Odor, TasteBonding TextureRadioactivity Malliability Ductile

23. 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

24. Scientists DemocritusFirst to believe matter made of atoms (atomos) AristotleMatter made of four elements (fire, air, earth and water) -Used by Church for nearly 1500 yrs. -Experiment to gain knowledge -Gases made of corpuscles (atoms) w/ spaces J. DaltonAtomic Theory -All elements made of atoms -All elements of same kind have same atoms -Reactions are changing, etc. of atoms -Conservation of Matter W. CrookesUsed CRT to find electrons JJ ThompsonDetermined charge on electron to be negative(magnets on CRT) -Plum Pudding Model of Atom R. MillikanMass and charge amount of electron -Oil Drop Experiment E. RutherfordGold Foil Experiment -Used alpha radiation (a He nucleus) to find protons -Beta and gamma radiation J. ChadwickFound neutron (no charge)

25. N. BohrAtoms give off certain colors/wavelengths of light -Planetary Model of Atom EinsteinBrownian Motion – atoms shown to exist by hitting pollen Photoelectric Effect – light is wave or photon Light (photon) is a packet of energy w/ no mass E. SchrodingerElectrons at speed of light Electron Cloud Model W. HeisenbergUncertainty Principle – never certain of election place Light used to see electrons makes them move P. DiracAnti-matter L. De BroglieElectron a wave and particle too Four Forces StrongHolds nucleus together with gluons WeakHold each neutron and proton together (quarks) w/ bosons EMFHolds nucleus to electrons - photons GravityNo one knows – maybe a graviton Standard Model of AtomIncludes subatomic particles (260+) Developed from Quantum Theory All matter made of Fermion particles Baryonic Matter made of quarks (Hadrons -3; Mesons – 2) Non-Baryonic includes leptons (electrons) QuarksMake up Baryons 3 in proton and 3 in neutron 2 in Mesons 6 Types – up, down, top, bottom, charm, strange

26. Four Universal Forces

27. Standard Model  Pauli Exclusion Principle – No two e - in the same atom can have the same quantum numbers In each suborbital, there will be two e -, and each will have opposite spins  Aufbau Principle – Atoms will fill up their orbitals starting with the lowest energies to highest

28.  Hund’s Rule – Each suborbital will have a maximum of 2 e - ; with opposite spins. Each suborbital in an energy level must have one e - before any receive a second one

29. Valency Valence Electrons – the electrons in the outermost shell or energy level – Octet Rule – atoms seek to fill to completion the outermost energy level by gaining or giving away electrons – Valence electrons will determine bonds and reactions – Valency provides grouping of elements on Periodic Table (Groups I through VIII)

30. Electron Configuration We need to be able to show how an atom’s electrons are configured (distributed) in order to explain chemical bonds and reactions There are several ways to do this, but first: – You must know and understand the element’s valence energy level (the period) – You must know the # valence e - (Group I to VIII or if a Transitional Metal) – And you must know the orbital grouping: s, p, d, or f

31. Aufbau Notation This gives the valence energy level, the valence orbital shape and the # of valence electrons in that level Example: 2 p 5 Energy Level Orbital # e -

32. Aufbau Notation: Long Format: – Na w/ 11 electrons would be: 1s 2 2s 2 2p 6 3s 1 – Short Format only gives the LAST Energy Level and the number of electrons in it Na  3s 1 – Without this e -, Na would isoelectric w/ Ne (have the same e - configuration) Cl -17 would be 1s 2 2s 2 2p 6 3s 2 3p 5 in Long Format Cl – 17 would be 3s 2 3p 5 in Short Format and its Lewis Dot Structure would have 7 dots around the elemental symbol

33. Often, you will see another shorthand way of writing an element’s electron configuration – This is called the Noble Gas Configuration – Noble Gases are in Column VIII – For example, using Na, Na’s configuration would be: [Ne]3s 1 – The [ ] means isoelectric (the starting point) – Li would be [He]2s 1 - La  [Xe]6s 2 5d 1 – Mn would be [Ar]4s 2 3d 5 - Ce  [Xe]6s 2 5d 1 4f 1 – Zn would be [Ar]4s 2 3d 10 - Pr  [Xe]6s 2 4f 3

34. Lewis Dot Diagrams Dot diagrams demonstrate the type of covalent bonds an element may make under certain conditions The element’s symbol is surrounded by up to 8 dots, each dot representing a valence e - Maximum of 2 dots per side (2 dots x 4 sides = 8 total, an octet valence level) It can be used for single atoms or to show molecules and their bonds and shapes

35. Determine dots from either doing the e - mapping OR learn to read the periodic table! Often, for transitional metals – need to do mapping (problem if ion) Remember Hund’s Rule Watch for + and – ions and either add or subtract the correct number of electrons Each side of elemental symbol represents one of the last 4 suborbitals (□) of the SAME Energy Level! Do not have to have 8 e - dots, can have openings – This is where bonds to other atoms can occur – Note: RN columns all have same dot drawing!

36. Lewis Dot Diagrams

37. Nucleosynthesis Due to the super high temperatures and density, the e -, e +, p +, and n 0 began to collide at the start of the universe – They got close enough for the Strong Force to take over allowing nuclei to form The first element was the nucleus of an isotope of Hydrogen, Deuterium ( 2 H) – a p +, n 0 This was joined by a p + and formed 3 He Then, with another n 0, 4 He formed (the stable form) With the capture of free e -, a stable atom was made Process called alpha combination

38. Nucleosynthesis in sum:

39. When dealing with Gases, there are four major variables utilized: Volume Pressure Temperature Amount of Gas (Moles)

40. Gas Equation Summary Boyle’s LawP 1 V 1 = P 2 V 2  If P goes up, V goes down (inverse) Charles’s LawV 1 = V 2 T 1 T 2  If V up, T up Gay-Lussac’s LawP 1 = P 2 T 1 T 2  If T up, P goes up Combined Gas LawP 1 V 1 = P 2 V 2 T 1 T 2

41. Dmitri Mendeleev (1834-1907) is given usually given credit for developing the first periodic table based on atomic weight and properties – it allowed him to predict new elements. His predictions were proven true with the discovery of gallium, scandium and germanium

42. However, it was Henry Moseley (1887-1915) who set it up using the atomic number (Z) He used x-rays to note the frequency of the element as well as the number of p + The chart became arranged according to atomic number rather than atomic weight

43. 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)

44. Z = atomic number = p + The number of protons identifies the atom and which element it is In a stable atom: – # p + = # n 0 = # e - – Thus, Na in its stable form has 11 p + ; 11 n 0 ; and 11 e - – If it has an unequal number of p + and n 0, then it is called an ISOTOPE

45. The Carbon Isotope

46. 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

47. 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

48. 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

49. Using the Periodic Table Elements in the Main Groups (A), form fairly consistent ions Group I will form +1 ions; Group II form up to +2; Group III form up to +3 ions Group IV will form either up to -4 or +4 ions Group V will form up to -3 ions; Group VI up to -2; Group VII form -1; and Group VIII will not form ions at all Those elements in the B group (transition metals) vary in their + charges meaning they can form different ions – look them up!

50. Group Names - Periodic Table

51. Know These!

52. Ions and Isotopes in Review Stable atom: #p + = #n 0 = #e - Atomic Mass: #n 0 = # 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

53. 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 n 0 (isotope) and e - (ion) change

54. Cf -3 has an atomic number of 98 – This means it has 98 p + – Its atomic mass is 216 – It has 118 n 0, (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 n 0 ; but due to the +1 charge, it has only 29 e -

55. 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

56. Counting Atoms in a Molecule In the example, NH 3, the subscript 3 only applies to the hydrogen. – Therefore: there is 1 N and 3 H in ammonia In the example, 3Ca 3 (PO 4 ) 2, the number of atoms changes due to the Coefficient is 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 3Ca 3 (PO 4 ) 2

57. 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

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

59. 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

60. Naming Covalent Compounds Process: 1.Prefix Indicating # + full name of first nonmetal 2.Prefix Indicating # + root name of second nonmetal + suffix “ide” 3.Watch for polyatomics and use their proper names

61. If it has more than two atoms – need to use the prefixes Number PrefixNumberPrefix 1 Mono7Hepta 2 Di8Octa 3 Tri9Nona 4 Tetra10Deca 5 Penta11Undeca 6 Hexa12Dodeca

62. For Example: NOTE THESE ARE ALL NONMETALS WITH NONMETALS! P 4 S 10 becomes Tetraphosphorous Decasulfide P 2 O 5 becomes Diphosphorous Pentaoxide SF 6 becomes Sulfur Hexafluoride N 2 O 3 is Dinitrogen Trioxide CO is Carbon Monoxide SO 2 is Sulfur Dioxide SiBr 4 becomes Silicon Tetrabromide Water is really Dihydrogen Monoxide!

63. Naming Ionic Compounds is really simple: 1. Name the cation (metal) using its proper name; if it is a polyatomic (NH 4 +1 ammonium), do the same 2. Then, using the stem of the anion (nonmetal mono/polyatomic), 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! [Iron (III) Sulfide] 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

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

65. NH 4 + + O -2 (NH 4 ) 2 O The parentheses are used to keep the polyatomic together; the 1 is understood Pb +4 + CO 3 -2 Pb 2 (CO 3 ) 4 and this can be simplified by reducing the subscripts to Pb(CO 3 ) 2

66. 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

67. Must remember which elements are normally diatomic (N 2, O 2, F 2, Cl 2, Br 2, I 2, and H 2 ) – AND S 8 and P 4 if by selves!!!!! 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

68. Example: H 2 + O 2  H 2 O – 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.....

69. 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.....

70. H 2 + O 2  H 2 O 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: 2 H 2 + O 2  2 H 2 O

71. 2H 2 + O 2  2H 2 O Now this is balanced! It means it takes 2 hydrogen molecules and one oxygen molecule to form 2 water molecules

72. 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

73. Let’s take the unbalanced equation of: KClO 3  KCl + O 2 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

74. 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 KClO 3  KCl + 3 O 2 This changes the number of K and Cl now You have to readjust again......

75. 2 KClO 3  KCl + 3 O 2 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 KClO 3  2KCl + 3O 2 K 2  K 2 Cl 2 Cl 2 O 6 O 6 It’s Balanced! Finally.

76. 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

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

78. DECOMPOSITION REACTIONS – Compounds break down into components AB  A + B or ABC  AB + C Examples... CaCO 3  CaO + CO 2 Ca(OH) 2  CaO + H 2 O 2 KClO 3  2 KCl + 3 O 2 2 NaCl  2 Na + Cl 2 H 2 CO 3  H 2 O + CO 2

79. 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 H 2 O  Metal OH - + H 2 (g) – Metals replace H in Acids  salt + H 2 (g) Al + H 2 SO 4  AlSO 4 + H 2 (g) 2 Sc(s) + 6 HCl (aq)  2 ScCl 2 (aq) + 3 H 2 (g)

80. DOUBLE REPLACEMENT Example: – FeCl 3 + 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

81. COMBUSTION – An exothermic rxn (gives off energy) – Usually find CO 2 and H 2 O in products – O 2 usually found in reactants CH 4 (g) + 2 O 2  CO 2 (g) + 2 H 2 O(g) + heat 2 C 4 H 10 (g) + 13 O 2 (g)  8 CO 2 (g) + 10 H 2 O(g)

82. ACID/BASE REACTIONS An acid + base  salt + H 2 O – 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

83. Exothermic Reaction – Gives off energy during reaction Endothermic Reaction – Absorbs heat/energy during reaction

84. Solutions (solns) Definition: – It is a homogeneous mixture – It is made up of a solute (what is being dissolved) and a solvent (what the solute is dissolved in) Solute + Solvent = Solution – It can be a gas, liquid, or a solid or even a combination of these phases – It does not separate into its parts on its own

85. Solubility If a solute dissolves in a solvent, it is soluble – Solubility affected by Temperature, Pressure, Agitation, etc. The process in which the solvent particles surround the solute and “cause” it to dissolve is called solvation – If this occurs in water – called hydration – Rule of Solvation – “Like Dissolves Like”

86. Rule of Solvation – “Like Dissolves Like” Refers to polar solutes dissolving in polar solvent, but not in nonpolar solvents This is due to the charges found on the molecules and opposite charges attracting each other If it does mix – said to be miscible, and if doesn’t – immiscible – Immiscible mixtures are not solutions since they will filter or separate on their own – Like oil and water

87. If the solute is an ionic compound and is dissolving completely, it is breaking down into its cations and anions – This is called ionization – Dependent upon the dissociation energy (D o ), the energy needed to break apart the ionic compound – This is a part of the solvation process

88. Saturation Saturation refers to the level of solute in the solvent If the solvent can not dissolve any more solute, then the two are in equilibrium and the solution is saturated If the solute “falls” out of the solvent – oversaturation If not enough solute - undersaturated

89. Solution Types Suspension - the solute will filter out of the solvent on its own (dirt in water) Thixotrope - the solute and solvent adt as a solid until agitated, then it acts as a liquid Emulsion - a liquid with a solid solute dissolved in it Aerosol - a colloid using a gas as a propellant Colloid - a heterogeneous mixture, the solute does not settle from solvent

90. Colligative Properties of Solns Properties of solns that depend on the number of solute particles in the solvent 1.Vapor Pressure – affecting evaporation rates. This is amount of solvent found in gas phase above the solution. Lowers vapor pressure. 2.Boiling Point – more solute, higher the B p 3.Freezing Point – more solute, lower F p 4.Osmotic Pressure – affects the passing through of semi-permeable membranes

91. Solutions Summary Soln = Solute (what is being dissolve) in a solvent (what it is being dissolved in) Solute Effects the Vapor Pressure (evaporation rate); the boil and freeze points and the osmotic pressure of the solution Saturation vs. Dilution – Precipitate Not easily filtered (based on particle size) – If mix – called miscible – “Like Dissolves Like” refers to polar and nonpolar substance mixing – Other Common Solns: Suspension – larger particle size means will filter out Emulsion – a solid solute in a liquid solvent Thixotrope – acts as a solid or liquid until agitated Colloid – hetereogeneous mixture where solute will not entirely filter out

92. Acid (pH) and Base (pOH) Strong Acid (1) Neutral (7) Strong Base (14)

93. A/B General Info Bases – Also called alkaline solutions – Have OH in them – React with acids in neutralization rxn to makes salts and water – Turn pink litmus paper blue – Taste bitter and feel soapy Acids – Usually have H in them with (aq) after their name – Ionize in water – React with bases to make salts and water – Turn blue litmus paper pink – Taste sour

94. A/B Rxn Acid (with H)+ Base (with OH)  Salt + H OH H w/ nonmetal Metal with OH  Metal (from B)Nonmetal(from A) + H 2 O Other Notes: – pH + pOH will equal 14 pH and pOH are INVERSE to each other pH measure ionization of H+; pOH measure ionization of OH- Dissociation of ions determines strength of A or B – As acid gets stronger, base weaker – an inverse relationship – There are 7 main Bases (most have OH in them) END OF CHEMISTRY REVIEW SECTION!!!

95. Start of Physics Review Section

96. Speed, Velocity and Acceleration The speed of an object is the distance the object travels per unit of time. Speed is a rate which tells you the amount of something that occurs or changes in one unit of time. Speed=distance over time Speed can be divided into two subtitles constant speed & average speed. Constant speed is the speed that does not change. Average speed is the total distance divided by time. Velocity is a speed in a given direction

97. Velocity V 1 represents the initial or starting velocity – If the object starts from a rest, V 1 will = 0 V 2 represents the final velocity of an object – If the object ends with a stop, then V 2 right at the end will be a zero, but not just a millisecond before that! – V = d / t – And this means d = vt; and t = d / v

98. Acceleration The acceleration of an object as produced by a net force is directly proportional to the magnitude of the net force, in the same direction as the net force, and inversely proportional to the mass of the object. Acceleration (a) = ΔV / Δt -or- Acceleration = force over mass

99. Newton's 1 st Law of Motion An object at rest tends to stay at rest and an object in motion tends to stay in motion with the same speed and in the same direction unless acted upon by an unbalanced force. unless acted upon by an unbalanced force Sometimes referred to as the “Law of Inertia." – Inertia is the state of rest or resisting a force that may cause motion or a change in velocity Frame of Reference – how the observer sees the change in velocity Frame of Reference (Point of View) can be stationary or moving depending on the observer Example: When a car stops suddenly, all the loose objects will continue forward until they hit something that stops them (have you ever had coffee do this at a stoplight?)

100. Newton's 2 nd Law of Motion The second law states that the acceleration of an object is dependent upon two variables - the net force acting upon the object and the mass of the object.net force It explains the relation of force, mass & acceleration. Force=mass x acceleration (F = ma) Weight is also a force = m x g The net force on an object is equal to the product of its acceleration and its mass.

101. Force Force is measured in the SI unit called a Newton (N) – 1 N = 1 kg x 1 m / s 2 1 N =.225 lbs 1 lb. = 4.448 N Forces usually are in equilibrium (balanced) Weight is a Force (wt = m g)

102. Force Continued... By definition it is a push or pull It can be divided into two subsets: unbalanced and balanced Unbalanced force can cause an object to start or stop moving; or change its acceleration, velocity or direction A balanced force is equal forces on an object that will not change the object’s motion

103. Acceleration - Due to Gravity a grav or just plain g, has a value of 9.80665 m/s 2 – We’ll round this off to 9.81 m/s 2 Use 10 for guesstimating! – Believe it or not – a grav at the equator is 9.7804 m/s 2 and at the poles it is 9.8321 m/s 2

104. Free Fall Acceleration If v 1 (initial velocity is zero or the object is at rest then falls): – V 2 = gt – V 2 = √2gh – H = ½ gt 2 – H = v 2 t 2

105. If v 1 does not equal zero... the object is thrown down or is shot downwards V 2 = v 1 + g t V 2 2 = v 1 2 + 2 g h H = v 1 t + ½ g t 2 H = v 2 + v 1 t 2

106. Momentum (ρ) Momentum is the product of an object’s mass and velocity It is directly proportional to mass and velocity It’s the tendency for an object to keep in motion – p = m v – F t = m v; where F t is the impulse or change in momentum

107. Newton’s 3 rd Law Basically – the law means that for every action there is an equal and opposite reaction A rocket launch – the F thrust downwards (action) forces the rocket upwards (reaction) against the F gravity Remember: Action  Equal/Opposite Reaction

108. Vectors Properties of Vectors – Vectors can be rearranged into a diagram – Size and Direction can not be changed – Use the tip to tail method of rearranging vector arrows – addition – To subtract vectors – add one to its opposite – Gives final displacement – Example: One Dimension A (5 m) + B (6 m) = R (Resultant) of 11 m to the right (same direction = addition) A (4m) + B (-3m) = R of 1 m to the left (opposite directions = subtraction)

109. Some Key Concepts Mass – the amount of matter something has Weight – mass affected by the force of gravity (m x g) – this a Force! Density – how much mass per volume (d = m / volume) Can be determined through math or through the displacement of fluid method *(Remember Archimedes and the crown)

110. Work, Power, Energy Work Work is a force applied to an object that causes displacement W = F Δ d – Measured in Newton-meters (Nm) or Joules (J) – Kg m 2 / s 2 is also called a Joule Power Power is the rate at which work is being done Measured in Joules per Second or Watts (W) and 1 J/s = 1 W Power (P) = Work / Time – P = W / t = F d / t Energy Potential Energy – The stored energy of position, inertia, or ability to do work – PE = m g h Kinetic Energy – Energy of motion – KE =.5 m v 2 – KE = F d

111. KE / PE Example As KE increases, PE decreases and versa vice As it moves upwards, height increases and PE increases; as it moves downwards, velocity increases and KE increases The pendulum at the two highest points have high PE, but no KE until it starts to move towards the center again. Then, the PE decreases until the bob hits the bottom and KE is at its highest

112. No medium needed All are transverse waves Have an electrical and magnetic field at right angles to each other Longest Wavelengths  Shortest λ Lowest f  Highest f Lowest energy (eV)  Highest eV Velocity is the same throughout = c = 300 000 km/s

113. Waves: f = Hz; λ = wavelength

114. Differences between Gravitational and Electromagnetic radiation There are two principal differences between gravity and electromagnetism, each with its own set of consequences for the nature and information content of its radiation, as described. Gravity is a weak force, but has only one sign of charge. Electromagnetism is much stronger, but comes in two opposing signs of charge. This is the most significant difference between gravity and electromagnetism, and is the main reason why we perceive these two phenomena so differently. Significant Gravitational fields are generated by accumulating bulk concentrations of matter. Electromagnetic fields are generated by slight imbalances caused by small (often microscopic) separations of charge. Gravitational charge is equivalent to inertia. Electromagnetic charge is unrelated to inertia.

115. SOUND Equations V sound = 331.5 +.61 (C o )If no temp given, assume 343 m/s v = d / t d = v t t = d / v  Denser the material, faster the sound! f = v / λIn Hertz (Hz) λ = v / f v = f λ v = λ / T (period) Intensity (I) = Power (P) / Area (A) – Intensity (I) = P / 4 π r 2 In Watts / meter 2 Power = I (4 π r 2 )In Watts Doppler Effect = f o = (v + v d / v + v s ) f s

116. Light, Mirrors and Lenses Convex MirrorConcave Mirror

118. Concave Mirror

119. Electricity and Magnetism

120. Magnitude of Charge Coulomb’s Law F Electric = K q q’ r 2 K = 8.988 x 10 9 Nm 2 /c 2 q and q’ are charges of objects r is distance between objects Coulomb (c) and Amperage (I) Amount of charge flowing through a wire in 1 second with a current of 1 ampere Ampere is 1 Coulomb per second, the intensity (I) of the electrical current Based on the charge of an electron 1 coulomb = 6.242 x 10 18 e - – Current (I) = Q / t in amperes Measuring the intensity of the electric current Charge of an electron (e-) = 1.60218 x 10 -19 c = 1 eV Potential Difference (V) Amount of work in an electric field to take the charge of 1 coulomb from one point to another Volt is the potential difference across a conductor that carries a current of 1 amp V = W / Q – V is potential difference in Volts One volt = J/c – W is work done in Joules – Q is charge in Coulombs

121. Resistance (R) Measured in Ohms  Ω R = V / I – I ohm (Ω) = 1 V / Amp Ohm’s Law  V = I R – Voltage = Current in Amps x Resistance in Ohms Resistance in Series – R 1 + R 2 + R 3 + …. = R Total Resistance in Parallel – 1/R 1 + 1/R 2 + 1/R 3 + …. = 1/R Total Capacitance (C) C = Q / V Measured in farads (1 coulomb per volt) Parallel Capacitance – C 1 + C 2 + C 3 + … = C Total Series Capacitance – 1/C 1 + 1/C 2 + 1/C 3 + … = 1/C Total Work and Power Work (W E ) = q V – In Joules Power (P) = V q / t – In Watts (J/s) – Power also = V I = I 2 R = V 2 / R

122. Magnetism Based on charges of atom’s particles It is a field force – line go from N to S (Faraday Lines) – Measured in Teslas or Gauss (1T = 100000G); – Earth =.0001T All magnets have two poles – if cut it makes new poles! Can lose magnetism if it is heated past material’s “Curie Temperature” and it returns when cooled Types of Magnetism: – Diamagnetic: no magnetism in material – Paramagnetic: magnetic only when in a magnetic field – Ferromagnetic: due to e- sea model of metal, it can be permanently magnetized