FINAL REVIEW —CHM 1032

The review has 2 parts: a partial outline for the course and practice questions to help you review the material and to help you study for the final. Your final is accumulative including everything you have been tested on, including chapter 11. 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 Law: summary of observations

o Cannot be violated o Allows you to predict future observations

Theory: unifying principle that explains facts and lawso Many laws can build a theory

States of Matter Fixed volume AND shape? __________________ Fixed volume and indefinite shape? __________________ Indefinite volume AND shape__________________

Classification of Matter Pure substances

o ___________________= 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 chemical composition

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

o Example: burning or bleaching

Measurement: **MEMORIZE

1mL= 1cc = 1 cm3

Significant Figures in Calculations: Multiplication/ Division: lowest # of sig figs Addition/ Subtraction: lowest # of sig decimals

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 elements

o Atoms combine in single, whole # ratios to form compoundso Atoms of one element cannot change into atoms of different atoms

Subatomic Particles: Atomic Number is:

Change in # of protons:

Change in # of neutrons: Change in # of electrons:

o Positive ions= o Negative ions=

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

Type of chemical compounds Ionic: Molecular:

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:

**MEMORIZE: POLYATOMICS and DIATOMICS

Calculations and Conversions

% Composition: Mass%=massof element X ¿compound ¿mass of total compound

×100%

% Yield: actual

theoretical×100%

Molarity= molesLiters

Dilution: M 1V 1=M 2V 2

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

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

o Molecular formula= empirical formula x n

o n= molarmass

empirical formulamass Structural: lines represent which atoms are bonded to each other

Types of Reactions Precipitation Reactions:

o Solubility Rules **Do not need to memorizeo Soluble with NO EXCEPTIONS= K+, Na+ and NH3

+

Acid-Base Reactions= Neutralizationo Acid: substance that produces H+

o Base: substance that produces OH-

Acid + Base -> salt + water Redox Reactions (OIL RIG)

o Oxidation is LOSING an electron

o Reduction is GAINING an electron

Representations of Chemical Equations-- Example

BaCl2 + MgSO4 o Molecular Equation:

o Complete Ionic Equation:

o Net Ionic Equation:

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 (+)

Quantifying heat—Calorimetry: measures heat of reaction at constant pressureo q= mCT

heat= mass x specific heat capacity x change in temperature uses specific heat: C with units: J/g* degrees Celsius

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

o v= cλ

o Greater wavelength= smaller frequency

Energy of a wavelength: hv=h×cλ

Plancks constant: h= 6.626 x10-34 J/s (--will be provided)***KNOW THE LIGHT SPECTRUM ORDER

Orbital Filling ***MEMORIZE order 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…

o Can use the periodic table for order in electron configuration 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 (“get on the bus”)

Aufbau principle: lower energy levels are filled firsto Orbitals can only hold 2 electrons in each lobe

Periodic Table Trends *** Memorize Trends

Lewis Structures – [1] count valence electrons [2] determine central atom –least electronegative except for Hydrogen [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

o When violating octets follow the same rules except you must keep count of valence in respect to the outer atoms

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)o Fill valence octet of outer atoms first, while maintaining the central atoms original

valence electron count

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- 1.9)o Ionic: 2+

*** Do not need to MEMORIZE geometry

Gases Pressure: atm is standard unit

o 1 atm= 760 mmHg (or torr) =101.325 kPa = 14.696 psi Temperature: Kelvin is standard unit Kelvin= C +273 **MEMORIZE Density= mass/ volume

o Can find mass from using PV= nRT to find moles, then convert STP= standard temperature and pressure

o Temperature= 273K o Pressure= 1 atm

Make sure to convert P, V, and T to correct units!

LAWS

Ideals gases we are assuming:o R= 0.0821 L*atm/mol*Ko The molecules have a very small/ negligible volumeo There are no interactions between the gas molecules

Dalton’s Law of Partial Pressure

Part 2:

1. Chemistry is a. The mystical search for the meaning of life.b. the science of what matter does by studying atoms and molecules.c. speculation about the nature of atoms and molecules.d. None of thesee. All of these

2. Which of these undergoes a chemical change?a. Melting iceb. Baking a cakec. Oxidation of copperd. B and Ce. All of the above

3. Which of the following is not a mixture?a. Gatoradeb. Oil in waterc. Aird. Iron nail

4. Determine the answer to the following equation with correct number of significant figures:(16.103 + 1.03) × 1.02521 = a. 17.5649b. 17.56c. 17.6d. 18

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

6. How many neutrons, protons, and electrons are in 15N?a. 8n, 6 p, 7eb. 7 n, 7 p, 8 ec. 7 n, 8 p, 7 ed. 8 n, 7 p, 7 e

7. Which of is not correct?a. O2

b. Cl2

c. C2

d. N2

e. H2

8. Which of the following is an ionic compound?a. C2H4O2

b. CaF2

c. H2Od. CCl4

9. What is the correct chemical formula when magnesium bonds with nitrogen?a. MnNb. MgNc. Mg2N3

d. Mg3N2

e. Mn3N2

10. Name each of the following: i. CCl4

ii. NaBr

iii. Ca(NO2)2

iv. H3PO4

11. What is the correct chemical formula for Copper (II) Phosphate?a. Cu3(PO4)2

b. CuPO4

c. Cu2(PO4)3

d. Co(PO3)3

12. Write out the formula for the following:1. Calcium Nitride

2. Sulfur Trioxide

3. Hydrofluoric Acid

4. Nitrous Acid

13. What is the correct name for the compound between nitrogen and fluorine?a. Nitrogen Fluorideb. Nitric Fluorinec. Fluorine Nitride

d. Mononitrogen Pentafluoridee. Nitrogen Pentafluoride

14. 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. Precipitation c. Combustion d. Decomposition

15. Finish the equation. What would precipitate when the reaction occurs?CuSO4 + NaOH

a. CuSO4

b. Cuc. Cu(OH)2

d. NaOHe. CuNa2

f. Na2SO4

16. In the following equation, which element is being reduced?I2 + CaO CaI2 + O2

a. Cab. Oc. I

17. Which of the following would be a nonmetal in period 4 of the periodic table?a. Siliconb. Galliumc. Carbond. Seleniume. Geranium

18. Select which of the following pairs are metals.a. Mg and Cab. Ge and Asc. Ag and Atd. B and Ce. All of the above

19. If an element had 34 protons, 34 neutrons, and 36 electrons what would form?a. -2 cationb. -2 anionc. +2 cationd. -2 anion

e. None of the above

20. A scientist found three isotopes of the same element X with natural abundances presented below. What is the atomic mass of element X?

Mass (amu) Abundance35.72 24%36.04 40%37.87 36%

a. 35.72b. 36.04c. 36.62d. 37.87

21. If an element in group 6A produced an ion, it would most likely have a charge of:a. 2+b. 2-c. 3-d. 3+e. 4-

22. What is the molar mass of C8H10N4O2?a) 200.16b) 180.22c) 168.15d) 194.22

23. Taking the molar mass of glucose from problem 21, how many moles would be in 125.87 g of glucose?

a) 0.629 molsb) 0.749 molsc) 0.648 molsd) 0.698 mols

24. Taking the moles from question 22, how many atoms of glucose are there?a) 4.51 x1023 moleculesb) 3.90 x 1023 moleculesc) 4.51 x 1023 moleculesd) 3.79 x 10 23 molecules

25. What is the mass ratio of oxygen to carbon in C3H9O?

26. Balance the equation : Ca(OH)2 + H3PO4 Ca3(PO4)2 +H2O

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

Equation: C2H6 + O2 CO2 + H2Oa. 20.3 gb. 18.5 gc. 16.7 gd. 12.3 g

28. How many moles of oxygen gas are in 5000 mL solution with a density of 1.429 g/L?a. 0.447 molb. 0.1366 molc. 0.223 mold. 0.137 mole. 0.2732 mol

29. If 53mL of nitric acid is mixed into .35L of water, what is the percent of acid?a. 13.2%b. 17.0 %c. 18.9 %d. 15.1 %

30. 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

31. How many moles are in a 198.6 mL solution with .523 molarity?a. 104 molb. .896 molc. 10.4 mold. .104 mole. 8.94mol

32. If you have a 100mL of a 4.5M solution, then dilute it until you have a 3L solution, what is the dilutions molarity?

a. 150Mb. 0.15Mc. 0.30Md. 3.00M

33. What is the empirical formula of C12H18O9 and what is the molar mass of the empirical formula of this compound?

a. C24H36O18: 612.6b. C6H9O4.5 ; 153.2c. CHO ; 29.0d. C4H6O3 ; 102.1

34. Out of the following options which one has the highest frequency?a. Radio Wavesb. Infraredc. Visible Lightd. Microwaves

35. What is the energy of a wave when it has a wavelength of 5.67 x 10 -9 m? 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

36. What is the electron configuration for Se?

37. What is the noble gas configuration of Bromine? a) [Ar] 1s22s22p64s23d104p5

b) [Ar] 4s23d104p5

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

38. What are the 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=2 b. n=2 → n=5 c. n=4 → n=3 d. n=3 → n=4 e. n=5 → n=6

40. Which of the following molecules is the most electronegative?a. Cb. Oc. Cad. N

41. Out of the bonds below, which of the following are the most polar?a. N------Brb. N------Sc. N------Fd. N------Cl

42. Draw the Lewis structure for CaF2

43. Draw the Lewis structures. Determine its # of electron groups and electron geometry:a. H2O

b. CCl4

c. SO3

44. Draw the Lewis structures. Determine its # of electron groups and molecular geometry:a. PF3

b. CO3 -1

c. SO2

45. Convert 76 degrees F to Kelvin:

46. 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 Ld) 13.4 L

47. 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?

e. .980 La) .765 Lb) 1.28 Lc) 1.34 L

48. 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

49. Which of the following statements are incorrect?a) As temperature increases, pressure increasesb) As temperature increases, volume increases

c) As pressure increases, volume increasesd) As pressure decreases, volume increases

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

A) True B) False

51. 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) 36200 mLc) 40100 mLd) 32400 mL