Chemistry SOL Review - Chem Site

Chemistry SOL Review - Chem Site

Chemistry SOL Review 1. Laboratory Safety Always wear goggles ! Never taste chemicals! To smell a chemical waft ! When mixing solutions ADD acid to water! Always rinse chemicals off skin with water !

Chemistry SOL ReviewScientific Investigation Safety What to you do if you spill anything on yourself in the lab? Identify three things that are unsafe in the picture below: 2. Lab Equipment Balance measures mass in ________ g, mg, kg ______ Beaker/Erlenmeyer Flask volume in

measures _______ mL, L _____ Graduated Cylinder volume measures _______ volume Pipet measures______ Crucible used for heating _______ 3. Scientific Method Parts of an Experiment

Independent variable: variable changed on purposegoes on x-axis Dependent variable: responding variablegoes on y-axis Control experiment: experiment where the independent variable is set to zero Constants: variables that are kept constant during a set of trials Analyze the following experiment and identify the control experiment, independent variable, dependent variable, and constants. A student designed this experiment to determine the effect of dissolving calcium chloride on water temperature. Different amounts of calcium chloride were added to room temperature water and the final temperature recorded. Trials

1 2 3 4 mL water 50 50

50 50 Starting water temperature 20C 20C 20C 20C

grams CaCl2 0 5 10 15 Final Water Temperature

20C 26C 31C 37C 4. Percent Error Used to tell how off you are from the value you should have gotten. Used mostly in lab. Ex: The specific heat capacity of iron is 0.45 J/gC. A student uses a calorimeter to experimentally determine the specific heat of

iron to be 0.60 J/gC. What is the students percent error? (Accepted experimental)/Accepted X 100 (0.45 0.60)/0.45 x 100 5. Graphing Indirect Relationship \ Direct Relationship / 6. Scientific Notation

Ex: 2.5 x 10-3 negative If the exponent is ________ then the smaller number in standard notation is _______ than 1 positive then the If the exponent is _______ greater number in standard notation is _______ than 1 7. Uncertainty and Significant

Figures When taking a measurement, always measure one decimal place past the scale of your instrument. For instance, the graduated cylinder to the left is measured with a 0.1 scale. The measurement recorded is 1.15 mL (1 place past the scale of the instrument). The 5 is the digit we are uncertain about. Significant Figures in Measurements: Non-zero digits are always significant. Any zeros between two significant digits are significant.

A final zero or trailing zeros in the decimal portion ONLY are significant How many significant figures does each number below contain? 3 123 ___ 3 103 ___ 1 0.001 ___ 5 10300. ___

3 10300 ___ 4 0.003010 ___ 8. Uncertainty and Significant Figures Addition and Subtraction The answer cannot have more places after the decimal than your measurement with the fewest places after the decimal. Ex: 2.59 + 2.3 = 2.9

4.506 cm + 2.9 cm = 7.406 7.4 cm 2.5 g - .36 g = 2.14 2.1 g Multiplication and Division The answer cannot have more significant figures than your measurement with the fewest number of significant figures. Ex: Ex: 2.500 x 2.0 = 5.0 6.5 x 3 = 19.5 20 100 / 4.00 = 25.00 30 9. Precision vs. Accuracy Precision - repeatability of results

Accuracy - getting the right answer 0.200 cm 0.190 cm 0.201 cm (accepted value = 0.201 cm) How would you describe these results? Accurate, but not precise 10. Temperature Conversions Celsius Kelvin K = C + 273

What is human body temperature in Celsius, Fahrenheit, and Kelvin? 392.98 K = _______ C = 119.88 11. Density D = mass/volume Units = g/mL, g/cm3 Density determines whether or not an object will: Sink or Float

l If an object has a mass of 5.0 g and a density of 20.0 g/ mL, what is the volume of the object? 20.0 g/mL = 5.0 g/V V = 0.25 mL A graduated cylinder is filled to the 10.0 mL line with water. A cube of tin (density = 7.3 g/mL) is placed in the graduated cylinder. The water level in the graduated cylinder rises to 20.0 mL. What is the mass of the cube of tin? 7.3 g/mL = m/10 mL M = 73 g 12. Metric Conversions

1000 mL = 1 L 1000 mm = 1 m 100 cm = 1m 1000 m = 1 km My house is 2.5 km from Deep Run. What is this distance in meters? 2.5 km 2500 m 13. Separating Mixtures

This figures shows an experimental setup used to separate solids form liquids. Which laboratory technique is shown on the right? A. Chromatograhy B. Filtration C. Decanting D. Distillation 14. Properties of States of Matter 15. Intermolecular Forces Intermolecular Attractions and Molecular Properties

As intermolecular forces increase, the molecules are held more strongly together. Solids resist melting because melting requires breaking intermolecular attractions and reforming new ones as the molecules slide past each other. Liquids resist boiling because the liquid molecules will have to overcome the intermolecular attraction of the other liquid molecules to enter the gas phase. 16. Chemical and Physical Changes Physical Changes: changes that do not affect the composition of the substance

Any change in the state of matter of a substance is a PHYSICAL change! Solid liquid = melting Liquid solid = freezing Liquid gas = evaporation Gas liquid = condensation Solid gas = sublimation 16. Chemical and Physical Changes Chemical Changes: changes in which a new substance is formed What are four signs that a chemical

reaction has occurred? Bubbles Color Change Heat Absorbed or Released Precipitate formed 17. Specific Heat Capacity Specific heat capacity: the amount of energy required to raise the temperature of 1 g of a substance by 1 degree Celsius low specific heat capacity, it heats up quickly.

If an object has a ____ high specific heat capacity, it heats up slowly. If an object has a ____ J/gC A 5.0 g object is heated from 25 C to 45 C. If it has a specific heat of 4.5 J/gC, what is the heat generated by the object? 18. Atomic Structure Using the SOL Periodic Table Lets use the periodic table to answer some questions about Silicon. How many protons does Silicon have? Atomic Structure

Using the SOL Periodic Table Lets use the periodic table to answer some questions about Silicon. How many protons does Silicon have? 14 protons = atomic number. How many electrons does neutral Silicon have? Atomic Structure Using the SOL Periodic Table Lets use the periodic table to answer some questions about Silicon. How many protons does Silicon have? 14 protons = atomic number. How many electrons does neutral Silicon have? 14 electrons (# electrons = # protons in neutral atoms) How many neutrons does Silicon-30 have? Atomic Structure

Using the SOL Periodic Table Lets use the periodic table to answer some questions about Silicon. How many protons does Silicon have? 14 protons = atomic number. How many electrons does neutral Silicon have? 14 electrons (# electrons = # protons in neutral atoms) How many neutrons does Silicon-30 have? 16 neutrons. Silicon-30 is an isotope of Silicon. It has a mass number of 30. The mass number is protons + neutrons. Atomic Structure Using the SOL Periodic Table Lets use the periodic table to answer some questions about Silicon. How many protons does Silicon have? 14 protons = atomic number. How many electrons does neutral Silicon have? 14 electrons (#

electrons = # protons in neutral atoms) How many neutrons does Silicon-30 have? 16 neutrons. Silicon-30 is an isotope of Silicon. It has a mass number of 30. The mass number is protons + neutrons. What is the molar mass of Silicon? Atomic Structure Using the SOL Periodic Table Lets use the periodic table to answer some questions about Silicon. How many protons does Silicon have? 14 protons = atomic number. How many electrons does neutral Silicon have? 14 electrons (# electrons = # protons in neutral atoms) How many neutrons does Silicon-30 have? 30 neutrons. Silicon-30 is an isotope of Silicon. It has a mass number of 30. The mass number is

protons + neutrons. What is the molar mass of Silicon? 28.0855 grams/mole (this is the same as the atomic mass on the periodic table) Atomic Structure Using the SOL Periodic Table Lets use the periodic table to answer some questions about Silicon. How many protons does Silicon have? 14 protons = atomic number. How many electrons does neutral Silicon have? 14 electrons (# electrons = # protons in neutral atoms) How many neutrons does Silicon-30 have? 30 neutrons. Silicon-30 is an isotope of Silicon. It has a mass number of 30. The mass number is protons + neutrons. What is the molar mass of Silicon? 28.0855 grams/mole (this is the

same as the atomic mass on the periodic table) How many valence electrons does Silicon have? Atomic Structure Using the SOL Periodic Table Lets use the periodic table to answer some questions about Silicon. How many protons does Silicon have? 14 protons = atomic number. How many electrons does neutral Silicon have? 14 electrons (# electrons = # protons in neutral atoms) How many neutrons does Silicon-30 have? 30 neutrons. Silicon-30 is an isotope of Silicon. It has a mass number of 30. The mass number is protons + neutrons. What is the molar mass of Silicon? 28.0855 grams/mole (this is the same as the atomic mass on the periodic table)

How many valence electrons does Silicon have? 4 valence electrons. Look for electrons in the highest principle energy level. 18. Atomic Structure Protons found in _______ nucleus have charge of ___ +1 Electrons found in electron ___________ __ cloud have charge of -1 Neutrons found in _______ nucleus have charge of __ 0

protons always equals the number The number of _______ electrons in a neutral atom. of ________ electrons In a magnesium ion, there are 2 more ________ protons giving the ion a total charge of +2. than ________ electrons than In a phosphide ion, there are 3 more ________ protons giving the ion a total charge of -3. ________ ELECTRONS CAN BE LOST OR GAINED!!! ONLY ____________

19. Isotopes/Ions/Atomic Structure Review Isotopes atoms of the same element neutrons with different numbers of ________. Ions - charged particles Symbol Atomic Number Atomic Mass

# protons # neutrons # electrons Charge P 15 31

15 16 15 0 Cl- 17 35

17 18 18 -1 Ca2+ 20 40

20 20 18 +2 Cl 17 37

17 20 17 0 37 17 20. Average Atomic Mass

The average atomic mass is an ________ average of all the isotopes of an element. (This is why the atomic mass on the periodic table is a decimal. That should make sense you cant have .01 neutrons!) Average Atomic Mass = (% abundance x mass number) + (% abundance + mass number) + There are two isotopes of chlorine, 35Cl which is 75% of the chlorine in the world, and 37Cl. What is the AAM of chlorine?

35.5 21. Scientists Thompson Model The atom is a positively charged diffuse mass with negatively charged electrons stuck in it. From Mark Rosengartens New York Regents Powerpoint 21. Scientists Rutherford Model

The atom is made of a small, dense, positively charged nucleus with electrons at a distance, the vast majority of the volume of the atom is empty space. Alpha particles shot at a thin sheet of gold foil: most go through (empty space). Some deflect or bounce off (small + charged nucleus). From Mark Rosengartens New York Regents Powerpoint 21. Scientists

Bohr Model Electrons orbit around the nucleus in energy levels (shells). Atomic bright-line spectra was the clue. From Mark Rosengartens New York Regents Powerpoint 21. Scientists Quantum-Mechanical Model Electron energy levels are wave functions. Electrons are found in orbitals, regions of space where an electron is most likely to be found. You cant know both where the electron is and where

it is going at the same time. Electrons buzz around the nucleus like gnats buzzing around your head. From Mark Rosengartens New York Regents Powerpoint 21. Chemists and their Contributions

Dalton: Atomic Theory / orbitals have 1 electron Pauli: 2 electrons per orbital / have opposite spin Heisenberg: Uncertainty Principle charge on electron Milikan: Democritus: coined word atom Hund: within a sublevel, dont pair e- until all neutrons Chadwick: Moseley: Periodic Table by Atomic # Mendeleev: Periodic Table by Atomic Mass

22. The orbitals and the periodic table 23. Electron Configurations Noble Gas Core Use noble gas before element as a shortcut s, p, d, f Blocks What is the electron configuration for Cd? [Kr]5s24d10 What is the configuration for the Cd2+ ion? [Kr]4d10

24. Orbital Diagrams Draw an orbital diagram for nickel 25. Family names Name the groups boxed in yellow, orange, green and blue. 26. Periodic Table Trends Atomic Radius

Ionization Energy Electronegativity Reactivity Period Group Down Up Up

Down Up Down Up Down Will Ca form an ion larger or smaller than the original atom? P? smaller, larger

27. Oxidation Numbers (Charges) loses or _____ gains an Charge results when an atom _____ ________. electron Metals _____ lose electrons, therefore become ________ positive ions cations called _______. negative Nonmetals _____ gain electrons, therefore become ________

called _______. anions -3 -2 +1 -1 +2 var 0 28. Valence Electrons Valence electrons electrons in the outer energy level (the

highest numbered energy level) 5 6 2 7 1 8 2 29. How do I tell if the Compound is Ionic or Covalent or Both?

Check to see what the compound is made up of: A metal and a nonmetalIts IONIC! 2 nonmetalsIts COVALENT! A polyatomic ion and another elementIts BOTH! (The polyatomic ion is the covalent part, the whole compound will be ionic.) 30. Ionic Bonds electrons Ionic bonds are formed when ________ transferred between a _____ metal and a

are _________ nonmetal ________. Non-metals above the staircase Metals below the staircase Non-metals above the staircase Metals below the

staircase The yellow shaded metals can take on multiple charges/oxidation states Types of Compounds (Ionic vs. Molecular) Ionic compounds form from metals and non-metals (across the tracks) and transfer electrons between elements. You figure out the formula for an ionic compound by criss-crossing charges to subscripts and reducing subscripts if possible. Ca2+ and F1- form ___________ Li1+ and PO43- form____________ Pb4+ and S2- form ________ Mn2+ and NO3-1 form _________

Types of Compounds (Ionic vs. Covalent) Ionic compounds form from metals and non-metals (across the tracks) and transfer electrons between elements. You figure out the formula for an ionic compound by criss-crossing charges to subscripts and reducing subscripts if possible. Ca2+ and F1- form CaF2 Li1+ and PO43- form Li3PO4 Pb4+ and S2- form Pb2S4 which reduces to PbS2 Mn2+ and NO3-1 form Mn(NO3)2 Naming Ionic Compounds

Write the name of the cation. If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion. If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses. Name the following compounds CaF2 _________________________ Li3PO4 __________________________ PbS2 _____________________________ Mn(NO3)2 ______________________________

Naming Ionic Compounds Write the name of the cation. If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion. If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses. Name the following compounds CaF2 Calcium fluoride Li3PO4 Lithium phosphate PbS2 Lead (IV) sulfide

Mn(NO3)2 Manganese (II) nitrate 31. Covalent Bonds Covalent bonds are formed when electrons are ______ shared between two ________ _________. nonmetals Covalent Compounds Covalent compounds are composed of two non-metals (above the staircase) Indicate # of each atom using prefixes (mono, di, tri, tetra, penta, hexa,

hepta, octa, nona, deca) The first element does not use mono if theres only one. Examples: OF2 is named oxygen diflouride N2O is named dinitrogen monoxide You try: NO2 ___________________________ P2O4 ____________________________ Molecular Compounds Molecular compounds are composed of two non-metals (above the staircase) Indicate # of each atom using prefixes (mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca)

The first element does not use mono if theres only one. Examples: OF2 is named oxygen diflouride N2O is named dinitrogen monoxide You try: NO2 nitrogen dioxide P2 O4 diphosphorus tetroxide 32. Polyatomic Ions

Nitrate Nitrite Sulfate Sulfite Phosphate Carbonate Hydroxide Ammonium

List formulas 33. Diatomic Elements hydrogen nitrogen

oxygen fluorine chlorine bromine iodine Remember! HNOFClBrI 34. Drawing Lewis Structures Dont forget Lewis Structures only use VALENCE Electrons! Draw structures for H2O, CO2, CCl4, and NH3

35. VSEPR Theory Valence Shell Electron Pair Repulsion Theory: basically means that the electrons want to be as far away from each other as possible Important shapes for the SOL: Shape Structure Example Bent

Draw H2O Trigonal planar Draw BF3 Trigonal pyramidal Draw

NH3 Tetrahedral Draw CH4 Linear draw CO2

36. Polarity Covalent bonds are when electrons are ______ shared nonmetals between two _________. If the electrons are shared equally, it is a ________ nonpolar covalent bond. If the electrons are shared unequally (meaning they are pulled closer to the more polar covalent electronegative element), it is a _____ bond.

36. Polarity To determine whether a bond is polar, nonpolar, or ionic, you must use a table of electronegativities. (This will be given to you on the SOL if you are supposed to use it.) When you subtract the two values, if the difference is between 0 and 0.4, the bond is nonpolar, meaning the electrons are shared equally between the two atoms between 0.4 and 2, the bond is polar, meaning the more electronegative element is pulling harder on the electrons greater than 2, the bond is ionic, meaning the more electronegative element pulled so hard on the electrons, that they came off one atom and were transferred to the other atom. 37. Writing Chemical Equations

REACTANTS PRODUCTS Write: Solid potassium chloride reacts with oxygen gas to yield solid potassium chlorate. KCl(s) + O2(g) KClO3(s) 37. Types of Chemical Reactions Synthesis: A + B AB Decomposition: AB A + B Single Replacement: AB + C AC +B Double Replacement: AB + CD AD + CB Combustion: CxHy + O2 CO2 + H2O Acid / Base: HX + MOH H2O + MX

38. Balancing Chemical Equations Balance equations to satisfy: the law of conservation of mass Write and balance: Magnesium reacts with nitrogen to yield magnesium nitride. 3Mg + N2 Mg3N2 39. Moles 1 mole = 6.022 x 1023 units 1 mole of gas at STP = 22.4 L

How many atoms are found in 10.0 g of sodium? 2.62 x 10 23 atoms 13 L of hydrogen at STP has a mass of 1.2 ___ g 40. Molar Mass grams / mole Also known as: Molecular weight Formula mass Formula weight Find the molar mass of potassium nitrate?

KNO3 = 101.11g 41. Percent Composition % composition = mass element / entire mass Find the percent magnesium in magnesium oxide? MgO 60 % 42. Stoichiometry ** Must have a balanced equation to solve these problems! Remember: grams to moles, mole ratio, moles to grams

2H2 + O2 2H2O How many grams of water will be produced from 5.0 g of hydrogen? 45 g H2O 43. Molecular and Empirical Formulas Molecular Formulas provide the true number of atoms in a compound Empirical formulas give the ratio of the elements found in a compound Structural formulas show how the atoms are connected. Molecular Formula Empirical Formula C6H6

CH C2H6 CH3 C2H2O4 CHO2 43. Empirical Formulas Empirical Formulas are the reduced form of Molecular formulas. For example: The empirical formula for C5H10 is CH2.

A favorite SOL type question: What is the empirical formula of a compound that contains 30% Nitrogen and 70% Oxygen? a) N2O b) NO2 c) N2O5 d) NO This is really a percent composition problem. Figure out which compound contains 30% nitrogen. 44. Kinetic Molecular Theory

The Major Points Temperature is related to kinetic energy Gas particles are in constant random motion Gas particles have no volume Kinetic Molecular Theory LIQUIDS When gas molecules lose kinetic energy (cool and slow down) then intermolecular forces can cause the molecules to stick together and liquify. Evaporation: molecules with enough kinetic energy to overcome the intermolecular attractions in a liquid can escape the liquid and enter the gas phase. Vapor Pressure: the force due to the gas above a liquid. This increases as temperature increases.

The curves are different for each liquid due to intermolecular forces Kinetic Molecular Theory LIQUIDS Boiling Point: the temperature where a liquids vapor pressure equals the external pressure or atmospheric pressure. Boiling Point increases as external/atmospheric pressure increases. Boiling Point decreases as external/atmospheric pressure decreases. Kinetic Molecular Theory LIQUIDS

Kinetic Molecular Theory SOLIDS 1. Particles in liquids are free to slide past each other 2. Particles in solids do not slide past each other, but vibrate in place. 3. Melting point: temperature where a solid becomes a liquid. 45. Gas Laws Boyles P1V1 = P2V2 @ constant temperature Charles's V1T2 = V2T1

45. Gas Laws Combined PV = PV T T 46. Ideal Gas Law PV = nRT Remember: No change occurs!

P = pressure in atm or kPa V = volume in L N = Moles R = constant (0.0821 L.atm/mol.K OR 8.314 L.kPa/mol.K) T = temperature in K 47. Endothermic Reactions Heat is ________. absorbed It appears on the ___ left side of the equation The quantity of heat will be positive ______.

48. Exothermic Reactions Heat is ________. released It appears on the ____ right side of the equation The quantity of heat will be negative _______. 49. Activation Energy The energy required to ____________. start a reaction A catalyst ______ lowers the activation energy.

50. Reaction Progress Diagram 51. Phase Diagrams critical point melting point boiling point 52. Heating Curves Temperature does not change during a phase change!

How much energy is required to melt 15.0 g of ice if the heat of fusion for water is 6.02 J/g? 90.3 J How much energy is required to raise the temperature of 15.0 g of water from 10 C to 25 C? 900 J (1 sig fig) 53. Kinetics Kinetics - Study of the rate of a reaction What are four things that affect the rate of a reaction?

Concentration Temperature Presence of catalyst Nature of reactants What is the collision theory? particles must collide for a reaction to occur 54. Catalysts Increase the rate of a reaction by:

lowering the activation energy Not used up in a reaction 55. Electrolytes An electrolyte dissociates _________ in a solution. (breaks up into ions) STRONG ELECTROLYTES: Conduct well Dissociate completely WEAK ELECTROLYTES

Conduct poorly Dissociate partially 56. Molarity Molarity = moles of solute/L of solution Calculate the molarity of a solution in which 15.0 g of NaCl is dissolved in 100. mL of water. 2.59 M 57. Dilution Molarity1 x Volume1 = Molarity2 x Volume2 What volume of a 4.0 M HCl solution should be used to make 100 mL of a 0.15 M HCl

solution? 0.00375 L (3.75 mL) 58. Solubility Curves How many grams of NaNO3 will dissolve in 100 g of water at 20 C? 85 g A supersaturated solution of KNO3 at 50 C would have more 85 than ___ g of solute in solution.

How many grams of KI will dissolve in 400 g of solution at 10 C? 540 g 59. Colligative Properties Properties that depend on how much solute is present Colligative Properties Adding impurities to a liquid increases the boiling point and decreases the freezing point (widens the liquid temperature range)

Examples: Adding antifreeze to the water in the radiator to prevent boiling in summer and freezing in winter. Putting salt on the road to prevent the road from icing up. lowers water's freezing point NaClNaCl decreases waters freezing point Freezing point, degrees C boiling point, degrees C NaCl increases water's boiling point

NaCl increases waters boiling point 130 125 120 115 110 105 100 0 100 200

300 grams of NaCl in 1Liter of water 400 0 -5 -10 -15 -20 -25 -30

0 100 200 300 grams NaCl in 1 L water 400 60. Chemical Equilibrium Equilibrium

when the concentration of reactants and products are constant Reversible reactions reactions that can go in either direction 61. LeChateliers Principle A reaction at equilibrium wants to stay at equilibrium. To accomplish this, the reaction will shift to the left or right to maintain equilibrium when a change is made. Shift Right

Shift Right 62. Acids Properties of Acids: H+ ions Low pH (can be negative) Tastes sour (vinegar) 63. Bases Properties of Bases: High pH OH- ions Bitter taste (soap, cleaning products)

Slippery Acid/Base Theory What is pH? pH indicates the hydrogen ion molarity [H+] in a solution pH = make [H+] exponent positive pOH indicates the hydroxide ion molarity [OH-] in a solution. pOH = make [OH-] exponent positive Example: A 1.0 x 10-3 molar solution of HCl would have a pH of ___ 3 4 Example: A 1.0 x 10-4 molar solution of KOH would have a pOH of ___ Memorize: pH + pOH = 14. 6

Example: A solution with a pH of 8 will have a pOH of: ____. 64. Titrations Add acid to base to find the molarity of either the acid or the base. An indicator changes color to show the endpoint of the titration. 65. Half Life A sample of element X has a half life of 8 days. If you start with 200 g of the sample, how much is left after 40 days? 6.25 g

66. Organic Chemistry Organic molecules have carbon. You cannot be asked anything specific to organic molecules, however you will most likely see organic molecules in other questions.

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