Final Study Guide Flashcards
To prepare for the final exam in Honors Chemistry. (75 cards)
1
Q
Hypothesis
A
A suggested solution
2
Q
Necessity for Experimental Controls
A
To keep the data pure/accurate.
3
Q
Qualitative
A
An observation that yields descriptive, nonnumerical results.
4
Q
Quantitative
A
Literal results, numerical etc. Yielding little rounded information, entirely factual.
5
Q
Mass vs. Weight
A
Mass is the whole of an object minus the factor of gravity- weight is mass + pull gravity
6
Q
SI unit of mass
A
Grams(g)
7
Q
SI unit of length
A
Meter(m)
8
Q
SI unit of Volume
A
Liter(l)
9
Q
Absolute Zero
A
K = C + 273.15
10
Q
Rules for Determining Significant Figures
A
- All non-zero digits are significant.
- All zeros between non-zero digits are significant.
- All beginning zeros are not significant.
- Ending zeros are significant if the decimal point is actually written in but not significant if the decimal point
is an understood decimal (the decimal point is not written in).
11
Q
Sig Fig Rules for Addition and Subtraction
A
The operation must yield a decimal that actuates to the same place as the smallest number of decimal places present.
12
Q
Sig Fig Rules for Multiplication and Division
A
The answer for a multiplication or division operation must have the same number of significant figures as the factor
with the least number of significant figures.
13
Q
Density Equation
A
Density = mass/volume
14
Q
Dimensional Analysis
A
A technique that involves the study of the dimensions (units) of physical quantities.
15
Q
Scientific Notation
A
The conservation of digits via e^a power, a means of condensing large numbers. Each reservation equates to one power in the e^to a power.
16
Q
Accuracy
A
how close a measurement is to the true value of the quantity being measured.
17
Q
Precision
A
refers to how close the values in a set of measurements are to one another.
18
Q
Percent Error
A
Percent error = accepted value - experimental value/ accepted value X 100%
19
Q
Matter
A
Anything that has mass and volume.
20
Q
Mixtures
A
Physical combinations of two or more substances.
21
Q
Pure Substance
A
Classified by the connection of elements and or compounds.
22
Q
Heterogeneous Mixture
A
Consists of visibly different substances
23
Q
Homogeneous Mixture
A
A mixture that appears to be one pure substance
24
Q
Element vs. a Compound
A
An element is comprised of a single form of matter ex: Li, Fe, Pb etc.
A compound is comprised of several elements that form matter
ex: H20, C02, NHCl etc.
25
Proper use of parentheses and subscripts in a chemical formula
The subscript serves on a individual basis else if there are multiple elements having a subscripts applied.
ex: Pb3(H20)4
26
How to Determine the Number of Atoms
Take the subscript value added (or multiplied in parenthetical cases) and calculate the elements based off the provided subscript relation
ex: Pb3 = 3 Pb's
H20 = 2 H's, and 1 O
27
Naming Compounds
1. Monatomic anions are named by replacing the end of the parent atom's name with "-ide."
2. The names of polyatomic ions do not change.
3. Ionic compounds are named by writing the name of the cation followed by the name of the anion.
28
Physical changes vs. Chemical Changes
Physical:
1. Changes of state (liquid to a gas, etc.)
2. Separation of a mixture
3. Physical deformation (cutting, denting, stretching)
4. Making solutions (special kinds of mixtures)
Chemical:
1. When one substance is changed into another
2. The formation of new chemicals.
3. Burning.
4. Change of color.
5. Formation of a precipitate.
29
Why is Democritus famed?
He is the father of the atom or "atmos".
30
Why is Dalton famed?
Dalton discovered the limitation in the law of definite proportions and established the law of definite proportions, and the law of multiple proportions.
31
Describe the law of definite proportions.
In a given type of chemical substance, the elements always combine in the same proportions by mass.
32
Describe the law of multiple proportions.
When two elements react to form more than one substance and the same amount of one element (like oxygen)
is used in each substance, then the ratio of the masses used of the other element (like nitrogen) will be in small
whole numbers.
33
Describe Thomson's plum-pudding model.
According to the
plum-pudding model, the negative electrons were like pieces of fruit and the positive material was like the batter or
the pudding.
34
Describe the Rutherford gold foil experiment.
One fires positively charged alpha particles, some bounce back. He (Rutherford) proposed that the positive matter was
concentrated in one spot, forming a small, positively charged particle somewhere in the center of the gold atom. We
now call this clump of positively charged mass the nucleus.
35
Identify the three major subatomic particles by charge, masses, and location in the atom.
Neutron: Located at the the center of the atom, has the greatest mass and is neutrally charged.
Proton: Located at the center of of the atom, has the second greatest mass and is positively charged.
Electron: Located around the outer regions of an atom, has the least mass and is negatively charged.
36
What is the definition of the Atomic Number?
An element’s atomic number (Z) is
| equal to the number of protons in the nuclei of any of its atoms
37
What is meant by the Atomic Mass of an element?
The atomic mass of an element is the average of the atoms present in an occuring sample of that element, described in "Amu's" aka "Atomic mass units".
38
Define the mass number of an atom.
The mass number is the total number of protons and neutrons in an atom's nucleus.
39
Describe isotopes.
Isotopes are atoms where the number of neutrons in the atom can vary.
40
Determine the number of protons, neutrons, and electrons in Li.
Lithium in the Table of Elements is indicated as having an Atomic Mass of 7, and, an Atomic Number of 3.
The number 3 indicates the number of Protons.
The Atomic Mass (7) is the number of Protons + Neutrons.
Therefore, total of 7 - 3 Protons = 4 Neutrons.
Now, As a Proton has a Positive Charge (+) and an Electron has a Negative charge (-), for the Atom to have no charge (Neutral) the the number of (+) must equal the number of (-).
Electrons therefore also equals 3. (2 in Energy Shell No. 1 and 1 in Shell No.2. This is the Valence Electron used in reactions with other Elements).
41
Boron has two naturally occurring isotopes. In a sample of boron, 20% of the atoms are B-10, which is an isotope
of boron with 5 neutrons and a mass of 10 amu. The other 80% of the atoms are B-11, which is an isotope of boron
with 6 neutrons and a mass of 11 amu. What is the atomic mass of boron?
Step One: Convert the percentages given in the question into their decimal forms by dividing each percentage by
100%:
Decimal form of 20% = 0:20
Decimal form of 80% = 0:80
Step Two: Multiply the mass of each isotope by its relative abundance (percentage) in decimal form:
20% of the mass of B-10 = 0:2010 amu = 2:0 amu
80% of the mass of B-11 = 0:8011 amu = 8:8 amu
Step Three: Find the total mass of the “average atom” by adding together the contributions from the different
isotopes:
Total mass of average atom = 2:0 amu+8:8 amu = 10:8 amu
The mass of an average boron atom, and thus boron’s atomic mass, is 10.8 amu.
42
What is the definition of the electromagnetic spectrum? What is the order of highest energy form to lowest energy form?
The electromagnetic
spectrum is the range of all possible frequencies of electromagnetic radiation. The highest energy form of
electromagnetic waves is gamma rays and the lowest energy form (that we have named) is radio waves.
43
Describe the appearance of an atomic emission spectrum.
Larger bandwidths equate to lower energy, whilst higher energy equates to smaller bandwidths.
44
Why can each element be identified by its emission spectrum?
An element can be identified by its emission spectrum because each emission is entirely unique to that element at hand.
45
Describe the Bohr electron cloud.
Each electron occupies a definite orbit that requires the electron to have a specific amount of energy.
46
Describe why the Bohr model of atoms explains the existence of atomic spectra.
Bohr was able to mathematically produce a set of energy levels for
the hydrogen atom. In his calculations, the differences between the energy levels were the exact same energies of
the frequencies of light emitted in the hydrogen spectrum.
47
What were the shortcomings of the Bohr Model?
1. It was determined that the energy levels in atoms with more than one electron could not be successfully calculated.
2. Another problem with Bohr’s theory was that the Bohr model did not explain why certain energy levels existed.
3. Yet another problem with the Bohr model was the predicted positions of the electrons in the electron cloud.
48
Name the model that replaced the Bohr model of the atom.
The electron cloud model.
49
Describe wave-particle duality.
Light travels as a wave and interacts with matter like a particle. Thus when light is
traveling through space, air, or other media, we speak of its wavelength and frequency, and when the light interacts
with matter, we switch to the characteristics of a particle (quantum).
50
Describe the relationship between the frequency and the energy of a photon.
f = E/h
51
Describe the concept of a standing wave.
Standing waves:
When a wave travels down a rope and encounters
an immovable boundary, the wave reflects off the boundary and travels back up the rope. This causes interference
between the wave traveling toward the tree and the reflected wave traveling back toward the person. If the person
moving the rope up and down adjusts the rhythm just right, the crests and troughs of the wave moving toward the
tree will coincide exactly with the crests and troughs of the reflected wave. When this occurs, the apparent horizontal
motion of the crests and troughs along the rope will cease.
52
State the Heisenberg uncertainty principle.
The principle states that it is impossible to know both the precise location
and the precise velocity of an electron at the same time.
53
State the relationship between the principal quantum number (n) , the number of orbitals, and the maximum number of electrons in a principal energy level.
Reference pg. 140 (look for the table at the bottom of the page)
54
Draw orbital representation for atoms.
Reference pg 149 (look for the picture at the top of the page)
55
Define valence electrons.
The electrons in the outermost shell.
56
Indicate the number of valence electrons for 4s23d104p1.
The 4s and 4p electrons can be lost in a chemical reaction, but the electrons in the filled 3d subshell cannot. Gallium
therefore has three valence electrons.
57
State the basis for the organization of Mendeleev's and Moseley's periodic table.
The table is arranged according to atomic number.
58
Identify the groups of the periodic table.
```
Alkali Metals
Alkali Earth Metals
Transition Metals
Boron Family
Carbon family
Nitrogen Family
Halogens
Noble Gases
Chalcogens
```
For details per group ex: 1A, 2A etc. (absolutely necessary!) visit pages 163 - 165.
59
How does one determine the number of valence electrons for a group in the periodic table?
Read the atomic number of the element.
60
Identify the purpose of periods in a periodic table?
The period corresponds to the energy level of that element's valence electrons.
61
Describe the similarities among elements in the same period in the periodic table.
The elements share the same number of orbital shells.
62
Explain the periodic law.
The periodic law states that the properties of the
| elements recur periodically as their atomic numbers increase.
63
Metals vs. Nonmetals vs Metalloids
Metals:
Ductile, malleable, conductive, an solid at room temperature (minus mercury).
Nonmetals:
Brittle, dull, poor conductors of heat and electricity, tend to gain electrons to form negative ions, and have low melting points.
Metalloids:
Have properties of both metals and nonmetals, often used as insulators, conductors, or semiconductors.
64
What is the stair-step line that separates the metallic elements from the nonmetallic ones?
That is the dividing line between metals and nonmetals.
65
Define atomic radius.
The atomic radius is one-half the
| distance between the centers of a homonuclear diatomic molecule.
66
Describe the factors that determine the trend in atomic size.
1. The number of protons (nuclear charge)
2. The number of energy levels
3. The shielding effect.
67
Describe the general trend in atomic size for groups and periods
Down increases, left increases. Up decreases, right decreases.
68
Describe the variations that occur in the general trend of atomic size in the transition metals.
This unusual electron configurations is due to that the d sub-level is particularly stable when it is half-full (5 electrons) or completely full (10 electrons)
69
Explain what an ion is.
An ion is an
| atom with a positive or negative charge.
70
Explain how cations are formed.
Cations are formed when an atoms looses electrons or gains protons.
71
Explain how anions are formed.
Anions are formed when an atom gains electrons or looses protons.
72
Explain why atoms form ions.
To achieve chemical stability.
73
Identify the atoms most likely to form positive ions.
Metals.
74
Identify the atoms most likely to form negative ions.
Nonmetals.
75
Determining Ionic Charge
MUST READ ALL OF 10.1 TO GRASP THIS CONCEPT!
