FINAL REVIEW—CHM 2045This is a partial outline for the course with some practice questions to help you review the material and in your studying for the final. Remember this a student made guide and may not have everything you need for the final. Good luck!

Matter: anything that has mass and occupies space Atom= submicroscopic particles, smallest objects that are called matter Molecules = two or more atoms held together by bonds

Scientific Approach to Knowledge Hypothesis: tentative interpretation or explanation for an observation

o A good hypothesis is ____________________________. Experiment: set of procedure designed to test hypothesis 2 Types of observations:

o Qualitative:___________________________o Quantitative: ____________________________

Law: summary of observationso Cannot be violatedo Allows you to predict future observations

Theory: unifying principle that explains facts and laws

States of Matter Fixed volume AND shape? ______________ Fixed volume and indefinite shape? ____________ Indefinite volume AND shape_____________

Classification of Matter

Pure substanceso __________________= made up of only one elemento __________________= made up of 2 or more elements

Mixtureso __________________= uniform throughouto __________________= not uniform throughout

Properties ___________________ change = is done without changing composition

o Example: change in state of matter (solid, liquid, gas) ___________________ change = chemical composition is changed

o Example: burning or bleaching

1. When H2O(g) H2O(s), what type of compositional change is it undergoing?

a) Condensationb) Sublimationc) Vaporizationd) Deposition

2. Which of these undergoes a chemical change?

a) Sublimation of dry ice into carbonb) Heating sugar for caramel c) Mixing an acid and based) B and Ce) All of the above

Measurement:

Accuracy: how close to the actual value a measurement gets Precision: how close the repeated measurements are to each other Significant Figures:

o Zero Rules: ______________________________o Calculations:

Multiplication/ Division: lowest # of sig figs Addition/ Subtraction: lowest # of decimal places

Atomic Laws Law of Conservation of Mass: mass is neither created nor destroyed Law of Definite Proportions: all samples of a compound have the same proportion of

their constituent elements, regardless of where the samples were taken Law of Multiple Proportions: when two elements (A and b) form two different

compounds, the masses of element B that combine with 1 gram of element A can be expressed as a ratio of small whole numbers

Atomic Theory: o Each element is composed of atomso All atoms of a given element have the same properties that distinguish them

from other elementso Atoms combine in single, whole # ratios to form compoundso Atoms of one elements cannot change into atoms of different atoms

Isotypes VS IonsChange in # of ____________ change in # of ___________

Positive ions: ______________ Negative ions: _______________

3. True or False: Protons and Neutrons contribute most of the mass of an atom.a) True B) False

4. How many neutrons, protons, and electrons are in 14C?

a) 8n, 6 p, 6eb) 6 n, 6 p, 6 ec) 6 n, 8 p, 6 ed) 6 n, 6 p, 8 e

Avagadro’s # = 6.022 x 1023 (a.k.a. 1 mole)

Elements that make up chemical bonds: Ionic: Covalent: Metallic:

Naming Different Compounds

Naming Ionic Compounds if metal only has 1 type of cation_________________+________________

Naming ionic compounds if metal has more than 1 type of cation________________(________________) + ________________

Naming Molecular Compounds

________________+_________________ ________________+________________ Naming Binary Acids

________________+__________________ __________________

Naming Oxyacids_________________ __________________

5. What is the correct molecular formula for Cobalt (III) Phosphate?

a) Co3(PO4)3

b) CoPO4

c) Co2(PO4)3

d) Co(PO3)3

6. Write out the formula for the following:

i. Lithium Nitride

ii. Sulfur Trioxide

iii. Hydroiodic Acid

iv. Potassium Sulfate

v. Nitrous Acid

vi. Silver (II) Fluoride:

7. What is the chemical formula for Magnesium Bromide?

a) MgBrb) Mg2Br2

c) MgBr2

d) Mg2Br

8. Name each of the following:

i. LiF

ii. FeCl2

iii. CsSO3

iv. HBr

v. NaNO3

vi. H2CO3

Calculations and Conversions

% Composition: Mass%of element X=massof element X∈1moleoof compoundmassof 1moleof compound

×100%

o Theoretical yield : calculated amount of product based on chemical equationo Actual yield: actual amount of product produced when doing reaction

% Yield: actual

theoretical×100%

o Limiting Reagent:___________________________________o Excess Reagent:____________________________________

Molarity= molesLiters

Dilution: M 1V 1=M 2V 2

Mass to Mass Calculations: Mass A-> Moles A-> Moles B-> Mass B

9. If glucoses chemical formula is C6H12O6 what is it’s molar mass?

a) 100.156b) 180.156c) 168.146d) 164.156

10. Taking the molar mass of glucose from problem 9, how many moles would be in 125.87 g of glucose?

a) 1.26 molsb) 0.749 molsc) 0.699 molsd) 0.767 mols

11. Taking the moles from question 10, how many atoms of glucose are there?

a) 4.21 x1023

b) 7.59 x 1023

c) 4.51 x 1023

d) 4.62 x 10 23

12. How many moles of nitrogen gas are in a 5000 mL solutions with a density of 0.765 g/L?a) 0.140 molb) 0.1366 molc) 0.273 mold) 0.137 mole) 0.2732 mol

13. What is the percent of oxygen in the compound phenol (C6H6O)?a) 16.8%b) 17.0 %c) 18.9 %d) 25.7 %

14. A chemical equation has 23.5 g of O2. How many grams of CO2 are produced in the equations?

Equation: C2H6 + O2 CO2 + H2O

A) 20.3 gB) 18.5 gC) 16.7 gD) 12.3 g

15. If 16.2 g of O2 react with 32.4 g of C2H5OH, what is the limiting reagent when we obtain 27.2 g of H2O?

a)C2H5OHb) O2

c) H2Od) none of them

16. How many moles are in a 98.4 mL solution with .678 molarity?

a) .0667 molb) .896 molc) .0954 mold) .0776 mol

17. A solution starts out with a molarity of .521 M and a volume of 79.6 L. If we are trying to dilute the solution and put it into 113.7 L, what would the new molarity be?

a) 3.65 x 10-4 Mb) 3.65 Mc) .365 Md) .0365 M

18. If there are 54.6 g of LiCl, in a 425 mL solution, what is the molarity of the solution?

a) 4.24 Mb) 2.78 Mc) 3.90 Md) 3.03 M

Formulas Empirical: lowest possible number of subscript Molecular: actual number of atoms in molecule

o Molecular formula= empirical formula x n

o n= molarmassempirical formulamass

Structural: lines represent which atoms are bonded to each other

19. What is the empirical formula of C6H18N9 and what is the molar mass of the empirical formula of this compound?

a) C6H18N9: 216.2b) C3H9N4.5 ; 108.1c) CHN ; 27.0d) C2H6N3 ; 72.11

Solubility Rules Soluble with NO EXCEPTIONS= K+, Na+, and NH3

+

Representations of Chemical Equations: Example: Fe(NO3)3 (aq) + NaOH (aq)

Molecular Equation: ________________________________________________________________________

Complete Ionic Equation: ________________________________________________________________________

Net Ionic Equation: ________________________________________________________________________

20. Balance the equation: HCN + NiSO4 H2SO4 + Ni(CN)2:

a) 2,1,1,1b) 2,1,2,1c) 1,1,2,2d) 2,1,1,2

21. Balance the equation: C2H4O2 + O2 CO2 + H2O

A) 1,6, 3, 4B) 1,1,2,2C) 1,2,2,2D) 1,2,2,1

Types of Reactions Acid-Base Reactions

o Acid: substance that produces H+

o Base: substance that produces OH-

Neutralizationo Acid + Base -> salt + watero Example: CH3COOH(aq) + NaOH(aq) ->CH3COONa(aq) + H2O(l)

Titration: used to quantify the concentration of an unknown acid or base using the opposite with a known concentration

Redox Reactionso Oxidation is LOSING an electron

o Reduction is GAINING an electron

o Rules to assigning oxidation numbers

Elements in their elemental form have an oxidation number of 0 The oxidation number of a monatomic ion is the same as its charge Nonmetals tend to have negative oxidation numbers, although some are

positive in certain compounds or ions Oxygen using has an oxidation number of -2 except in peroxides, when it

then becomes -1 Ex: H2O vs H2O2

Hydrogen is -1 when bonded to a metal, and +1 when bonded to a nonmetal

Ex: H2O vs AlH2

Fluorine always has an oxidation of -1 The sum of the oxidation numbers in a compound is equivalent to its

overall charge Ex: NO3

- N=+5; O= -2 -> (+5) + 3(-2) =-1

22. Which compound in the equation form the precipitate when the reaction occurs?LiBr + HgSO4 HgBr2 + Li2SO4

a) LiBrb) HgSO4

c) HgBr2

d) Li2SO4

23. From the equation above, what would make a compound a spectator ion?

a) Needs to be a solid at the endb) Needs to be on the reactants sidec) Needs to be on the products sided) Needs to be aqueous

24. Strong acids when mixed with strong bases form a salt and water as their products. What is the name of this type of reaction?

a) Neutralization

b) Precipitationc) Combustiond) Decomposition

25. Strong acids are used in the body to break down compounds for digestion. Which of the following is classified as a strong acid?

a) NaOHb) HClc) CH3COOHd) H2C2O4

26. In the following equation, which compound is the oxidizing agent, and is being reduced?Cl2 + NaBr NaCl + Br2

a) Clb) Nac) Br

Energy Generally measured in Joules Energy conversion Factors

o 1 calorie =4.184 jouleso 1 Calorie = 1000 calorieso 1 kilowatt- hour= 3.60 x 106 joules

Exothermic (-) vs Endothermic (+)

Kinetic Energy: KE=1

2mv2

Relationship between Internal Energy (E), Heat (q), and work (w)o E= q+w

Enthalpy: sum of internal energies and the product of the pressure and volumeo H=E +PV

o H= E+ PV

Quantifying heato q= mCT

heat= mass x specific heat capacity x change in temperature - qsystem= qsurroundings

Calorimetryo Coffee cup: measure heat of reaction at a constant pressure

qsystem= msystemCsTo Bomb: measures internal energy of a reaction at constant volume

q =CcalT Hess’s Law

o If chemical equation multiplied by some factor, the H is also multiplied by that factor

o If chemical equation is reversed, H changes signso If chemical equation can be expressed as a sum of a series of steps, H for

overall reaction is the sum of H for each step.

Standard enthalpies of formation (use with Lewis Structures)o Hreaction= n Hf(products) - n Hf(reactants)

27. Which of the following is an example of kinetic energy?

a) A child at the top of a slideb) A molecule when hitting the wallc) A bowling pin about to be hit by a bowling balld) A quarter back running for a touch down

28. Which of the following is an example of an exothermic reaction?

a) Combining to elements togetherb) A candle flamec) Baking muffinsd) Melting ice

29. (True/False): Work has a positive value when work is being done by the system.

a)True B) False

30. When 1 mol of fuel is used to light a gas grill it gives off 4512 J of heat, and does 37 J of work what is the ∆H and ∆E of the equations?

a) +4512 J, +4549 Jb) -37 J, +4549 Jc) -4512 J, -4549 Jd) +4549 J, -4549 J

31. A 25.6 g of copper initially at 79.8 C, is put in a container with 30.9 g of water initially at a temperature of 89.0 C. what is the final temperature of the container when the two items are placed together?

a) 89.7 Cb) 112.5 Cc) 65.8 Cd) 104.5 C

32. What is the ∆Hrxn of the equation, and is this equation endothermic or exothermic?C12H22O11 + O2 CO2 + H2O

∆H Values:C12H22O11= -2226.1 O2=0CO2=-393.5H20= -483.6

a) -1349, exothermicb) -3103.2, exothermicc) +1349, endothermicd) +3103.2, endothermice) None of the above

33. Calculate ∆ H rxn for the reaction: CaO (s) + CO2 (g) → CaCO3 (s)

Use the following ∆ H ’s.Ca (s) + CO2 (g) + 1/2O2 (g) → CaCO3 (s) ∆ H=−812.8 kJ2Ca (s) + O2 (g) → 2CaO (s) ∆ H=−1269.8 kJ

Electromagnetic Radiation speed of light (c)= 3.00 x108 m/s Frequency= # of waves that pass a point during a period of time

o v= cλ

o Greater wavelength= smaller frequency

Energy of a wavelength: E=hv=h×cλ

De Broglie Equation:

34. What is the energy of a wave when it has a wavelength of 5.67 x 10 -9 m? ( Hint: speed of light and planks constant are in the equation)

a) 3.51 x 10 -17 Jb) 2.67 x 10 -17 Jc) 5.87 x 10-17 Jd) 1.12 x 10 -17 J

35. Out of the following options which one has the lowest frequency?

a) Radio Wavesb) Infraredc) Visible Lightd) Microwaves

Energy of a Photon: E=hv/molesQuantum Numbers

n= principal energy level ( the 1 in 1s) l= angular quantum number

o s= 0o p= 1o d= 2o f= 3

m (ml)=magnetic quantum numbero can be between –l and +l

o exact value cannot be individually placed s= quantum spin; can be + or –

Orbital Filling Electrons occupy the lowest energy orbital available (1s before 2s) Fill in the following order -> 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s…

Pauli exclusion principle: no 2 electrons in an atom can have the same four quantum

numbers Hund’s Rule: when filling degenerate orbitals, electrons filled them singly first, with

parallel spins Aufbau principle: lower energy levels are filled first Orbitals can only hold 2 electrons each

Electron Configuration Writing all of the energy levels that are filled with electrons ( and how many in each) Short hand= using nearest noble gas preceding the element in question to shorten the

electron configurationo Li= [He] 2s1

o Noble gas ______________ be its own abbreviation Neon’s electron configuration cannot be [Ne]

36. What is the electron configuration for Hg?

37. What is the noble gas configuration of Selenium?

a) [Ar] 1s22s22p64s23d104p4

b) [Ar] 4s23d104p4

c) [Kr]d) [Ar] 3d104p6

38. How many valences electrons are in the electron configuration for Magnesium?

a) 2p6

b) 2s2

c) 1s2

d) 3s2

39. Which of the following transitions results in absorption of the of the highest-energy photon?

a) n=6 → n=2b) n=2 → n=5c) n=4 → n=3d) n=3 → n=4e) n=5 → n=6

40. Which of the following transitions results in the shortest wavelength of light? a) n=6 → n=2b) n=2 → n=5c) n=4 → n=3d) n=3 → n=4e) n=5 → n=6

Periodic Table Trends

Electron Affinity

Atomic Radius

Ionic Radius Trends for ions: Cations are _______________ than parents Anions are _______________ than parents

Ionization Energy

Electronegativity

Effective Nuclear Charge Zeff= Z- S(core electrons)

41. Which of the following molecules is the most electronegative?

a) Sb) Oc) Brd) N

42. Element C has the first three following Ionization energies, 610, 810, and 4500 KJ/mole. What is most likely the formula?a) XPO4

b) X2PO4

c) X2(PO4)3

d) X3(PO4)2 e) X(PO4)3

Characteristics of Elements

Alkaline Metals Alkaline Earth Metals Halogens Noble GasesLeftmost groupForms +1 cation

Group 2AForms 2+ cation

Group 7AForms 1- anion

Rightmost groupHas its octet (does not form ions)

43. Select which element is a metalloid.

a) Mg and Cab) Ge and Asc) Au and Hgd) B and Ce) All of the above

44. Metals form ions with:

a) noble gasesb) transition metalsc) with other metalsd) anions

45. Transition Metals are located in groups 3-12 on the periodic table, but have unique characteristics. What makes these transition metals unique compared to other elements?

a) They have a set charge and can bond to noble gasesb) They have set charge and all have same characteristics as one anotherc) They all can have different set of positive charges and bond to anionsd) They all have different set positive charges and can bond to each other

Lattice Energy: “Coulomb’s Law”

E=q1q2r

Lattice energy becomes less exothermic as ionic radius increases Lattice energy becomes more exothermic as magnitude of ionic charge increases

Lewis Structures [1] count valence electrons [2] determine central atom [3] distribute bonds between atoms, then add lone pairs to terminal atoms [4] any remaining electrons, after terminal atoms have their octet, go to the central

atom

Drawing for Ions:o Cations lose electrons

Become just the element’s symbol with the chargeo Anions gain electrons

Drawn as element’s symbol surrounded by valence electrons, [], and its charge

o Some can have incomplete octets (BF3) or expanded octet (SF6)

Bond Lengths Distance between nuclei of bonded atoms Single bond> double bond> triple bond

o Single bond= sigmao Double bond= sigma and pio Triple bond= sigma and 2 pi

The shorter the length, the stronger the bond

46. (True/False): Single bonds take the most energy to break.a) True B) False

Polarity Means that one atom has a stronger pull on the electrons than the other atom Differences in electronegativities to determine polarity

o Non polar: 0 – 0.4o Polar: 0.5- 1.9o Ionic: 2+

47. Out of the bonds below, which of the following are the most polar?

a) C------Ob) C------Nc) C------Fd) C------H

Hybridization [1] Count up sigma bonds and lone electron pairs [2] Determine the principal energy level of atom’s valence electrons [3] Take s-orbital to hybridize [4] Take as many p-orbitals and d-orbitals (if applicable) needed for same number of

orbitals hybridized as sigma bonds/ lone pairs

48. What is the electron, and molecular geometry of the SF4?

a) Trigonal bipyramidal, seesawb) Trigonal Bipyramidal, Trigonal Pyramidalc) Tetrahedral, Trigonal Pyramidald) Octahedral, square

49. What is the Orbital Hybridization of the following compound of the C attached to C in the compound C2H2?

a) sp3

b) spc) sp2

d) sp0

50. Luis structure time! This time add molecular geometry, electron geometry, number of sigma and pi bonds, indicate polarity, and hybridizations.

i. BF3

ii. HCN

iii. NH2-

iv. H2PO4 -

v. C2O4 -2

vi. XeO4

vii. SF4

viii. IF4

VSPER Theory/ Electron and Molecular Geometries

Gases Pressure: atm is standard unit

o 1 atm= 760 mmHg (or torr)o Manometer measures difference in atmospheric pressure and the pressure of

the gas in the vessel

Temperature: Kelvin is standard unito Kelvin= C +273

Density= mass/ volumeo Can find mass from using PV= nRT to find moles, then convert

STP= standard temperature and pressureo Temperature= 273Ko Pressure= 1 atm

51. Convert 85 degrees F to Celsius and Kelvin:

Laws Avagadro’s

oV 1

n1=V 2

n2o If pressure and temperature are constant

Boyle’so P1V1= P2V2

o If moles and temperature are constant Charles’

oV 1

T 1=V 2

T2o If moles and pressure are constant

Gay=Lussac’s

oP1T1

=P2T 2

o If volume and moles are constant Combined Gas

oP1V 1

T 1=P2V 2

T 2 Ideal Gas

o PV=nRTo R= 0.0821 L*atm/mol*Ko Make sure to convert P, V, and T to correct units!

Dalton’s Law of Partial Pressure

Diffusion vs Effusion Diffusion= collection of molecules spreading from high concentration to low

concentration Effusion= collection of molecules escaping through a tiny hole in a vacuum

o rate Arate B

=√ Molarmassof gas BMolar massof gas A

Van der Waals Equation –for real gases

52. A volume of liquid is 14.3 L with a pressure of 67.8 kPa, and is compressed to a pressure of 72.3 kPa. What is the new volume?

a) 10.9 Lb) 11.3 Lc) 17.8 L

d) 13.4 L

53. Your boiling potato’s in 1.2 L of water at 45 C on the stove top for 15 minutes. You check your potato’s 5 minutes later, and decide to turn up the heat to 67 C. What would the volume in the pot be when you turn up the temperature?

a) .980 Lb) .765 Lc) 1.28 Ld) 1.34 L

54. There is 5.6 moles of N2 in a container with a pressure at 465 kPa at a temperature 345 K. When the temperature is increased to 409 K, what is the new pressure due to the particles hitting each other at a faster rate?

a) 500 kPab) 551 kPac) 545 kPad) 567 kPa

55. When dealing with the combined gas law, which of the following statements are incorrect?

a) As temperature increases, pressure increasesb) As temperature increases, volume increasesc) As pressure increases, volume increasesd) As pressure decreases, volume increases

56. (True/False): At STP, pressure is 1 atm, and temperature is 25 C.

A) True B) False

57. What is the volume when a container has 32.4 g of H2O at a temperature of 56 C, and an internal pressure of 1.34 atm?

a) 40670 mLb) 36300 mLc) 40100 mLd) 32400 mL

58. A gas diffuses at a rate that is 1.63 compared to oxygen gas (constant temperature). What would be the Molar Mass of the unknown gas, and what element is the molar mass closest to on the periodic table?

a) 10.8, Bb) 11.3, C

c) 12.03, Cd) 13.8, N