Chemistry Midterm Ch 1-6 Review

Absolute zero
Accuracy
Analytical chemistry
Atom
Atomic emission spectrum
Atomic mass
Atomic mass unit
Atomic number
Atomic orbital
Atomic radius
aufbau principle
biochemistry
cation
Celsius temperature scale
chemical reaction
compound
distillation
electron
electron configuration
electronegativity
element
frequency
ground state
group
Heisenberg uncertainty principle
heterogeneous mixture
homogeneous 
inorganic chemistry
ionization energy
isotope
Kelvin temperature scale
mass
mass number
metal
mixture
neutron
nucleus
organic chemistry
Pauli exclusion principle
period
periodic law
phase
photon
physical chemistry
precision
product
proton
quantum
reactant
significant figure
spectrum
transition metal
vapor
wavelength
weight

1.	What is the temperature of absolute zero measured in C?
2.	The study of chemicals that, in general, do not contain carbon is traditionally called what type of chemistry?
3.	Which step in the scientific method requires you to use your senses to obtain information?
4.	The variable that is observed during an experiment is called what type of variable?
5.	The step that usually comes last in solving numeric problems is ____.
6.	How do conceptual problems differ from numeric problems?
7.	An example of a homogeneous mixture is ____.
8.	A substance that can be separated into two or more substances only by a chemical change is a(n) ____.
9.	What is one difference between a mixture and a compound?
10.	What do chemical symbols and formulas represent, respectively?
11.	A chemical change occurs when a piece of wood ____.
12.	What is the result of multiplying 2.5  10 by 3.5  10?
13.	What is the result of adding 2.5  10 and 3.5  10?
14.	Which group of measurements is the most precise? (Each group of measurements is for a different object.)
15.	Three different people weigh a standard mass of 2.00 g on the same balance. Each person obtains a reading of 7.32 g for the mass of the standard. These results imply that the balance that was used is ____.
16.	What is the measurement 111.009 mm rounded off to four significant digits?
17.	When multiplying and dividing measured quantities, the number of significant figures in the result should be equal to the number of significant figures in ____.
18.	The weight of an object ____.
19.	A cubic meter is about the same as the volume occupied by a ____.
20.	What is the density of an object having a mass of 8.0 g and a volume of 25 cm?
21.	What is the volume of 80.0 g of ether if the density of ether is 0.70 g/mL?
22.	As the density of a substance increases, the volume of a given mass of that substance ____.
23.	Who was the man who lived from 460B.C.–370B.C. and was among the first to suggest the idea of atoms?
24.	Dalton's atomic theory included which idea?
25.	Why did J. J. Thomson reason that electrons must be a part of the atoms of all elements?
26.	Who conducted experiments to determine the quantity of charge carried by an electron?
27.	What is the relative mass of an electron?
28.	As a consequence of the discovery of the nucleus by Rutherford, which model of the atom is thought to be true?
29.	The nucleus of an atom is ____.
30.	The mass number of an element is equal to ____.
31.	Using the periodic table, determine the number of neutrons in O.
32.	How many protons, electrons, and neutrons does an atom with atomic number 50 and mass number 125 contain?
33.	Select the correct symbol for an atom of tritium.
34.	How is the number of neutrons in the nucleus of an atom calculated?
35.	Why do chemists use relative masses of atoms compared to a reference isotope rather than the actual masses of the atoms?
36.	The atomic mass of an element depends upon the ____.
37.	In Bohr's model of the atom, where are the electrons and protons located?
38.	The principal quantum number indicates what property of an electron?
39.	What is the shape of the 3p atomic orbital?
40.	What is the maximum number of f orbitals in any single energy level in an atom?
41.	What is the maximum number of d orbitals in a principal energy level?
42.	What is the maximum number of orbitals in the p sublevel?
43.	The shape (not the size) of an electron cloud is determined by the electron's ____.
44.	The letter "p" in the symbol 4p indicates the ____.
45.	What types of atomic orbitals are in the third principal energy level?
46.	What is the next atomic orbital in the series 1s, 2s, 2p, 3s, 3p?
47.	According to the aufbau principle, ____.
48.	What is the number of electrons in the outermost energy level of an oxygen atom?
49.	What is the electron configuration of potassium?
50.	What is the basis for exceptions to the aufbau diagram?
51.	How does the speed of visible light compare with the speed of gamma rays, when both speeds are measured in a vacuum?
52.	Which color of visible light has the shortest wavelength?
53.	How are the frequency and wavelength of light related?
54.	As changes in energy levels of electrons increase, the frequencies of atomic line spectra they emit ____.
55.	The atomic emission spectra of a sodium atom on Earth and of a sodium atom in the sun would be ____.
56.	What is the approximate energy of a photon having a frequency of 4  10 Hz? (h = 6.6  10 Js)
57.	What is the approximate frequency of a photon having an energy 5  10 J? (h = 6.6  10 Js)
58.	Who predicted that all matter can behave as waves as well as particles?
59.	Each period in the periodic table corresponds to ____.
60.	Of the elements Pt, V, Li, and Kr, which is a nonmetal?
61.	The atomic number of an element is the total number of which particles in the nucleus?
62.	What element has the electron configuration 1s2s2p3s3p?
63.	Elements that are characterized by the filling of p orbitals are classified as ____.
64.	Which subatomic particle plays the greatest part in determining the properties of an element? 
65.	What are the Group 1A and Group 7A elements examples of?
66.	Atomic size generally ____.
67.	The metals in Groups 1A, 2A, and 3A ____.
68.	Why is the second ionization energy greater than the first ionization energy?
69.	For Group 2A metals, which electron is the most difficult to remove?
70.	Compared with the electronegativities of the elements on the left side of a period, the electronegativities of the elements on the right side of the same period tend to be ____.
71.	What is the sum of 6.210 L and 3 L expressed in the correct number of significant digits?
72.	The mass of the electron is 9.109 39  10 kg. Express the mass of the electron to 1, 2, 3, and 4 significant figures.
73.	What is the temperature 128 K expressed in degrees Celsius?
74.	Chlorine has two naturally occurring isotopes, Cl-35 and Cl-37. The atomic mass of chlorine is 35.45. Which of these two isotopes of chlorine is more abundant?
75.	Consider an element Z that has two naturally occurring isotopes with the following percent abundances: the isotope with a mass number of 19.0 is 55.0% abundant; the isotope with a mass number of 21.0 is 45.0% abundant. What is the average atomic mass for element Z?
76.	What is the electron configuration of sulfur?
77.	In which group in the periodic table do the elements have the highest electronegativity values?
78.	What is the atomic number for an element with 41 neutrons and a mass number of 80?
79.	How many electrons are in an atom of gold?
80.	How many neutrons are present in an atom of the isotope U?
81.	How many electrons are in the highest occupied energy level of copper?
82.	How many electrons are there in the highest occupied energy level of atoms in Group 5A elements?
83.	How many electrons are present in the d sublevel of a neutral atom of nickel?
84.	What is the usual charge on an ion from Group 7A?
85.	How did Lavoisier transform the field of chemistry in the late 1700s?
86.	List and describe the two steps for solving conceptual problems.
87.	Explain what is meant by a substance and a pure substance. Give an example of each.
88.	Why are you unable to observe the wavelike motion of a soccer ball as it is kicked toward a goal?
89.	Describe the trends in the atomic size of elements within groups and across periods in the periodic table. Provide examples.