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. Define physical and chemical change. Physical changes are state changes.
Chemical changes result in new substances.
3. Find the density of objects. What are the units for density? density=
mass/volume g/mL kg/L g/cm3
4. 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
5. Qualitative vs. quantitative measurements qualitative- do not deal with
quantity, merely that characteristics of a substance quantitative- depends on
how much
6. Differentiate between precision and accuracy. Precision- measurements are
close to one another accuracy- measurements are close to actual value
7. 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
8. Independent and dependent variables. How do you find the slope of a
line?Dependent variables are measured against the ind. Variable. The
independent variable stands alone. Slope= y2-y1
X2-x1
9. 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.
10. What are the branches of chemistry? Define each. Organic- study of carbon
Inorganic- study of all elements Analytical- precise measurements Physicalphysics of chem Theoretical- computers Biochem- chemistry of living things
11. 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
12. Who discovered the electron? How? Thomson- cathode ray tube
13. 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.
14. Who discovered the proton and nucleus? How? Rutherford- gold foil
experiment
15. Who discovered the neutron? Chadwick
16. Differentiate between mass and weight. Mass is the amt of matter present.
Weight depends on gravity.
17. Compare and contrast emission and absorption spectra. Emission spectra
show the wavelengths of light being emitted. Absorption spectra shows color
of what is being absorbed…black lines are what is being emitted.
18. 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
19. What is the equation that relates energy to frequency of light? E= h
20. What is Planck’s constant? h= 6.626 x 10-34 Js
21. What does Heisenberg’s uncertainty principle say? The position and velocity
of an electron cannot be determined simultaneously.
22. 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
23. 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.
24. 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.
25. Who is credited with developing periodic law? Mendeleev
26. Who organized the modern periodic table? How? Moseley, used increasing
atomic number
27. What do families on the periodic table have in common? Have same number
of valence electrons and same charge as ions
28. 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
29. 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.
30. 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.
Electron affinity- ability of an element to attract electrons
31. 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 EA- difficult to say
32. What can be determined by finding the electronegativity difference between
two elements? The type of bond that will form between them.
33. 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
34. 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
35. What are the differences between polar and nonpolar bonds? Polar bonds
have electrons shared unevenly. Nonpolar have electrons shared evenly.
36. Draw the Lewis structures for the following: SO4-2, SF6, NH3, NO3-1, BeCl2, CH4,
PCl5, H2O, HCN See board.
37. What is resonance? Resonance exists when more than 1 Lewis structure is
necessary to describe the true structure of a molecule.
38. What are the geometries according to the VSEPR theory and bond angles of
the molecules above (#36)? See board.
39. 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.
40. 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.
41. Which elements are always diatomic molecules in nature? HOFBrINCl
42. 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? Ca(OH)2
43. 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) C14H18N2O5 empirical and molecular.
44. What is the percent composition of iron in Fe2O3? 70%
45. 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.
46. 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.
47. What are the general types of chemical reactions that we discussed?
Synthesis, combustion, decomposition, single replacement, double
replacement, redox, acid-base
48. What forms when metal carbonates decompose? Metal oxide and carbon
dioxide
49. What forms when a metal reacts with an acid? Hydrogen gas and the ions of
the metal and counter ion
50. What are the products when metal chlorates decompose? Metal chloride plus
oxygen gas
51. What products are formed when hydrogen peroxide decomposes? Water and
oxygen gas
52. What are the products of a hydrocarbon or alcohol combustion? Water and
carbon dioxide
BE AWARE THAT YOU NEED TO KNOW NOMENCLATURE AND POLYATOMIC
IONS!