Name ______________________________________ Date _______________________ Period __________
Final Exam Study Guide
Pre-AP Chemistry
1. Differentiate between pure substance, mixture, elements, compounds,
molecules and ions. A pure substance has an identical composition and can
only be broken down by chemical means. A mixture can be separated by
physical means. Elements are examples of pure substances. Compounds are
also pure substances, and consist of a combination of elements held together
by ionic bonds. Molecules are a combination of elements held together by
covalent bonds. Ions are atoms that have gained or lost electrons.
2. Find the density of a block that has 3.0 cm sides and a mass of 4.5 g. What
are the units for density? 4.5 g/3.0x3.0x3.0= 0.5g/cm3
3. What are the 4 factors that indicate a chemical change will happen?1. color
change 2. Formation of a gas 3. Formation of a precipitate 4. Absorbing or
releasing heat
4. Define qualitative and quantitative measurements. qualitative- do not deal
with quantity, merely that characteristics of a substance quantitativedepends on how much
5. Differentiate between precision and accuracy. Precision- measurements are
close to one another accuracy- measurements are close to actual value
6. Be able to convert numbers in and out of scientific notation. a) 6.03 x 105
603,000b) 5.09 x 10-2 .0509c) 2305 2.305 x 103 d) 0.000986 9.86 x 10-4
7. Where are independent and dependent variables located on a graph? How
do you find the slope of a line? Dependent variables (y-axis) are measured
against the ind. Variable (x-axis). The independent variable stands alone.
Slope= y2-y1/x2-x1
8. What are the rules for determining significant figures? 1. All nonzero
numbers are sig. 2. Captive zeros are sig. 3. Leading zeros are NEVER sig.
4. Trailing zeros are only sig if a decimal is in the number.
9. What are isotopes? How can you recognize that you are dealing with an
isotope? Atoms with different masses but same atomic numbers. Mass
number will be different than weight on per. table
10. Who discovered the electron? How? Thomson- cathode ray tube
11. Describe the model of the atom proposed by Thomson. Rutherford? All
positive and negative charges were evenly dispersed. Plum pudding model
Rutherford’s model had a positive nucleus surrounded by negative electrons.
12. Who discovered the proton and nucleus? How? Rutherford- gold foil
experiment
13. Who discovered the neutron? Chadwick
14. Differentiate between mass and weight. Mass is the amt of matter present.
Weight depends on gravity.
15. Calculate the molar mass of the following: CuSO4, FeCl2, V2O5
CuSO4=159.62 g/mol, FeCl2= 126.75 g/mol V2O5= 181.88 g/mol
16. How many moles are in 3.4 g of Zn? 3.4g/65.38 g/mol= .052 mol
17. How many atoms are present in 4.5 moles of I? 2.7x1024 atoms
18. Calculate the number of moles in 23.6 g of Na3PO4. 23.6g/163.94 g/mol =
.144 moles
19. What is the molar mass of I2? How many atoms are present?MM= 253.8
g/mol atoms= 2 moles x 6.022 x 1023 = 1.20 x 1024
20. Why is emission spectra important to the discovery of electrons? The
emission spectra is unique for each element due to the number and
transitions of electrons.
21. Draw and describe the photoelectric effect. The photoelectric effect occurs
when a light of a minimum frequency is shown on a metal surface and an
electron is ejected.
See drawing on scanned page
22. Define the following: quanta, frequency, wavelength, speed of light,
amplitude. Small packet of energy, waves past a point in 1 second, distance
along a wave from two equal successive points, 3 x 108 m/s, the height of a
wave
23. What is the equation that relates energy to frequency of light? E= h
24. What is Planck’s constant? h= 6.626 x 10-34 Js
25. Calculate the wavelength of light with a frequency of 3.4 x 1015 Hz. c/=
3.0x108/3.4 x 1015= 8.8 x 10-8 m
26. What energy does light with a wavelength of 680 nm have? c/=
3.0x108/680 x 10-9 m=4.4 x 1014 Hz THEN E=h E= (6.626x10-34)( 4.4 x
1014 Hz )= 2.9x10-19 J
27. What does Heisenberg’s uncertainty principle say? The position and velocity
of an electron cannot be determined simultaneously.
28. What are Hund’s rule, Aufbau principle and Pauli exclusion principle? Hund’s
rule- electrons must be put in each equal energy orbital first before being
paired. Aufbau principle- electrons must occupy the lowest energy level
available first Pauli Exclusion principle- electrons in equal energy orbitals
must have opposite spins
29. How is an electron configuration for an element different from its orbital
notation? the orbital notation is more specific showing exactly which orbital
each electron goes in.
30. Write the electron configuration for the following: Zn, Xe, Ca, Br
Zn= 1s22s22p63s23p64s23d10
Xe= 1s22s22p63s23p64s23d104p65s24d105p6
Ca= 1s22s22p63s23p64s2
Br= 1s22s22p63s23p64s23d104p5
31. Draw orbital notation for: C, Ne, Sr
32. How are atoms converted into ions? What are cations and anions? Atoms
gain or lose electrons to form ions. Cations are positive ions that have lost
electrons. Anions are negative ions that have gained electrons.
33. Who is credited with developing periodic law? Mendeleev
34. Who organized the modern periodic table? How? Moseley, used increasing
atomic number
35. What do families on the periodic table have in common? Have same number
of valence electrons and same charge as ions
36. List all the families that we discussed on the periodic table and their group
numbers. Group 1- Alkali metals Group 2- Alkaline Earth metals Group 13Boron group Group 14- Carbon group Group 15- Nitrogen group Group 16Chalcogens Group 17- Halogens Group 18- Noble gases
37. In what state of matter do each of the halogens exist in nature? F, Cl are
gases. Br is liquid. I and At are solids.
38. Define the following: ionic radius, electronegativity, ionization energy, atomic
radius, electron affinity. Ionic radius- distance from nucleus to outer electron
cloud Electronegativity- The ability to attract electrons in a chemical bond.
Ionization energy- the energy required to remove an electron Atomic radiushalf the distance between the nuclei of two identical atoms bonded together.
39. What are the trends of the above on the periodic table? IR- increase down,
decrease across until group 14 then increases EN- increase across, decrease
down IE- increase across, decrease down AR- increase down, decrease
across
40. Which has the largest radius: Ca, N, or Xe? Xe
41. What can be determined by finding the electronegativity difference between
two elements? The type of bond that will form between them.
42. Define ionic, covalent and metallic bonds. Ionic- between metal and
nonmetal, give/receive electrons Covalent- between 2 nonmetals, share
electrons Metallic- between 2 metals, d orbitals overlap and electrons roam
freely
43. List two physical properties of ionic bonds, covalent bonds and metallic
bonds. Ionic bonds- high melting points, poor conductivity as solids Covalent
bonds- low melting points, poor conductivity Metallic bonds- high melting
points and good conductivity
44. What are the differences between polar and nonpolar bonds? Polar bonds
have electrons shared unevenly. Nonpolar have electrons shared evenly.
45. Draw the Lewis structures for the following: SO4-2, SF6, NH3, NO3-1, BeCl2, CH4,
PCl5, H2O, HCN
See answer on scanned page
46. What is resonance? Resonance exists when more than 1 Lewis structure is
necessary to describe the true structure of a molecule.
47. What are the geometries according to the VSEPR theory and bond angles of
the molecules above (#45)? Tetrahedral, octahedral, trigonal pyramidal,
trigonal planar, linear, tetrahedral, trigonal bipyramidal, bent, linear
48. What is a sigma bond? A pi bond? A sigma bond refers to s orbital overlap.
There is 1 sigma bond in single, double and triple bonds. Pi bonds refer to p
orbital overlap. There is 1 pi bond in a double bond and 2 pi bonds in a triple
bond.
49. How can you determine the hybridization of a central atom in a molecule?
Look at the number of bonds plus lone pairs on the central atom. Then count
them and use the s, p d and f notation.
50. Differentiate between single, double and triple bonds. Single bonds are the
longest and easiest to break. Double bonds are shorter and stronger. Triple
bonds are shorter and the strongest.
51. Which elements are always diatomic molecules in nature? HOFBrINCl
52. What is the law of conservation of mass? How is it important to balancing
equations? The law of conservation of mass says that mass cannot be created
or destroyed in a chemical reaction. It is important to balancing equations
because it means that you must have the same number of atoms on each side
of a balanced equation.
53. What are the general types of chemical reactions that we discussed?
Synthesis, combustion, decomposition, single replacement, double
replacement, redox, acid-base
54. What forms when metal carbonates decompose? Metal oxide and carbon
dioxide
55. What forms when a metal reacts with an acid? Hydrogen gas and the ions of
the metal and counter ion of the acid
56. What are the products when metal chlorates decompose? Metal chloride plus
oxygen gas
57. What products are formed when hydrogen peroxide decomposes? Water and
oxygen gas
58. What are the products of a hydrocarbon or alcohol combustion? Water and
carbon dioxide
59. What factors affect the rate of a reaction? 1) temperature 2) concentration 3)
surface area 4) catalysts
60. How can you tell if a reaction is endothermic or exothermic??? If the products
have less energy than the reactants, it is exothermic. If the products have
more energy than the reactants, it is endothermic.
61. Define LeChatelier’s principle. LeChatelier’s principle says that when a stress
is applied to a system at equilibrium, the system will shift to return to
equilibrium.
62. What does it mean when a reaction reaches equilibrium? The rates of the
forward and reverse reaction are equal
63. Find the empirical formula of a substance with 27.4% potassium, 34.7%
manganese and 40.5% oxygen.
See answer on scanned page
64. A compound was analyzed and found to contain 13.5 g Ca, 10.8 g O, and 0.675
g H. What is the empirical formula of the compound?
See answer on scanned page
65. NutraSweet is 57.14% C, 6.16% H, 9.52% N, and 27.18% O. Calculate the
empirical formula of NutraSweet and find the molecular formula. (The molar
mass of NutraSweet is 294.30 g/mol)
See answer on scanned page
BE AWARE THAT YOU NEED TO KNOW NOMENCLATURE AND POLYATOMIC
IONS!
Bonus!!!!
66. What are the ways that you can convert to moles… from grams, from volume at
STP, from atoms? Grams to moles, divide by molar mass. Moles to grams,
multiply by molar mass. 22.4 L= 1 mole of gas at STP Atoms to moles, divide
by avogadro’s number. Moles to atoms, multiply by avogadro’s number.
Draw the mole map!