Chemistry Semester 2 Final Exam Flashcards
(123 cards)
1
Q
Specific patterns in the periodic table that make it possible to make predictions about an element’s properties
A
periodic trends
2
Q
As the atomic number increases, the nuclear attraction for electrons increases, thus pulling the electrons in more tightly to the nucleus
A
Coulomb’s law
3
Q
One-half the distance between the nuclei of two identical atoms bonded together
A
atomic radius
4
Q
A measure of how readily an atom will react
A
reactivity
5
Q
When electrons in the inner energy levels block the attraction of the nucleus to the valence electrons
A
shielding effect
6
Q
A measure of the ability of an atom in a bond to attract electrons
A
electronegativity
7
Q
it is a measure of a neutral atom’s ability to gain an
electron
A
electron affinity
8
Q
The energy required to remove one electron from a neutral atom of an element
A
ionization energy
9
Q
the charge of an ion
A
oxidation numbers
10
Q
The radius of a monatomic ion in an ionic compound
A
ionic radius
11
Q
A force/mutual electrical attraction that holds atoms together in a substance
A
chemical bond
12
Q
A measure of the ability of an atom in a bond to attract electrons
A
electronegativity
13
Q
The tendency of atoms to prefer to have 8 valence electrons
A
Octet rule
14
Q
Compound formed when 2 or more elements are covalently bonded
A
molecule
15
Q
A binary ionic compound between 1 kind of metal atom and 1 kind of nonmetal atom
A
salt
16
Q
Network of cations and anions that are mutually attracted to one another
A
crystal lattice
17
Q
A compound made of positive and negative ions that chemically bond in a way that the charges equal out.
A
ionic compound
18
Q
charged atoms
A
ion
19
Q
Positively charged atom from losing electrons
A
cation
20
Q
Negatively charged atom from gaining electrons
A
anion
21
Q
A positively or negatively charged, covalently bonded group of atoms
A
polyatomic ion
22
Q
Metals that “transition” because they can form many different ions
A
transition metal
23
Q
An interaction that holds metal atoms together through a shared pool of valence electrons
A
metallic bond
24
Q
Electrons not associated with a single atom or bond
A
delocalized electron
25
A solution of two metals mixed together
alloy
26
A compound formed when two or more atoms share electrons
covalent compound
27
The bond between nonmetallic atoms that share electrons
covalent bond
28
A neutral group of atoms held together by covalent bonds
molecule
29
When 2 atoms of the same element exist naturally as a molecule
diatomic elements
30
Process by which substances collide with enough energy that new bonds are formed between atoms, thus creating new substances
chemical reaction
31
starting substances
reactants
32
ending substances
products
33
dissolved in water
aqueous
34
A binary compound with at least 1 atom of oxygen
oxide
35
A compound made of cations and anions
salt
36
The decomposition of a substance by an electric current
electrolysis
37
A compound made of carbon and hydrogen
hydrocarbon
38
An insoluble solid that forms from the ions of two aqueous compounds
precipitate
39
A list of elements in the order in which they will easily undergo certain chemical reactions
activity series
40
Summarize the connection between Coulomb’s law and the periodic trends.
↑ atomic number, ↑ the number of p+
↑ number of p+, ↑ the strength of positive nuclear charge
↑ strength of (+) nuclear charge, ↑ attractive force on e-
This law explains why we see the trends we see in atomic radii, reactivity, electronegativity, etc.
It is all about how tightly the electrons are being held to the nucleus, which is impacted by the # of p+ in the nucleus, and the number of e- 🡪 both the total and specifically in the outermost energy level.
41
Ionic radii:
Atoms that lose e- become smaller
Atoms that gain e- become larger
Tends to increase as you go DOWN a group
42
Explain how the shielding effect works.
This occurs when electrons in the inner energy levels block the attraction of the nucleus to the valence e-, and thus they become more loosely held.
It increases as you go down a group as the atomic size is increasing.
43
Describe how the difference in electronegativity between two atoms can be used to predict the type of bond that could form. Include an example to support your description.
ΔEN shows if one atom will more strongly attract the e- than the other. If the attraction is super strong (ΔEN > 2), an ionic bond forms because one takes the e- from the other. If the attraction is very low (ΔEN < 0.5), a covalent bond forms and the e- are mostly equally shared. If ΔEN is somewhere in the middle, a polar covalent bond is formed where e- are unequally shared.
Ex. Atmospheric nitrogen (N2) makes a ΔEN = 0, thus a covalent bond will form and e- will be equally shared.
44
Identify the oxidation numbers for elements in groups 1, 2, and 13-17.
Group 1 🡪 +1
Group 2 🡪 +2
Group 13 🡪 +3
Group 14 🡪 +/- 4
Group 15 🡪 -3
Group 16 🡪 -2
Group 17 🡪 -1
45
Differentiate between the forces between different substances versus the forces within a substance.
Intermolecular forces are between different substances
Intramolecular forces are within the same substances (chemical bonds)
46
Explain what is meant by the statement, “Bonding is a spectrum.”
There aren’t just two types of bonds. It is a spectrum with ionic and nonpolar covalent at the two extremes.
47
Write a sentence to summarize the connections between chemical bonds, chemical reactions, and compounds.
Chemical bonds form through chemical reactions and result in new compounds.
48
Explain why most atoms form chemical bonds and why some atoms don’t.
They form bonds to be stable with full outer energy levels of electrons.
Noble gases don’t form chemical bonds because they are already stable.
49
Explain the difference between what a chemical formula tells you about an ionic versus a covalent compound.
In a covalent compound, the chemical formula tells you exactly the number and types of atoms.
In an ionic compound, it represents the ratio of cations to anions in the crystal lattice.
50
IONIC OR COVALENT: electrons are transferred
ionic
51
IONIC OR COVALENT: between metal and nonmetal
ionic
52
IONIC OR COVALENT: always between ions
ionic
53
IONIC OR COVALENT: can result in the formation of salt
ionic
54
IONIC OR COVALENT: electrons are shared
covalent
55
IONIC OR COVALENT: always between nonmetals
covalent
56
IONIC OR COVALENT: can be polar or nonpolar
covalent
57
IONIC OR COVALENT: results in the formation of a molecule
covalent
58
IONIC OR COVALENT: crystalline solids
ionic
59
IONIC OR COVALENT: high melting and boiling points
ionic
60
IONIC OR COVALENT: can conduct electricity when dissolved in water
ionic
61
IONIC OR COVALENT: can be solid, liquid, or gas
covalent
62
IONIC OR COVALENT: low melting and boiling points
covalent
63
IONIC OR COVALENT: cannot conduct electricity when dissolved in water
covalent
64
List the number of atoms of each element in the compounds below. AlF3
Aluminum - 1
Fluorine - 3
65
K3P
Potassium - 3
Phosphorous - 1
66
Ca(ClO3)2
Calcium - 1
Chlorine - 2
Oxygen - 6
67
Explain the connection between electronegativity and the formation of an ionic bond.
The difference in electronegativity is so great that one atom takes the electrons from the other because it has such a greater attraction to them.
68
Use an example of an ionic compound to explain the “rule of zero charge”.
MgCl2 is one atom of magnesium with a 2+ charge, and two atoms of chlorine, each with a 1- charge.
1(2+) + 2(1-) = 2+ + 2- = 0
69
Differentiate between binary ionic compounds, compounds with polyatomic ions, and binary molecular compounds.
Binary ionic compounds: made of 2 elements that transfer e- in an ionic bond
Compounds with polyatomic ions: made of one (or more) covalently bonded charged group of atoms that transfer e- in an ionic bond.
Binary molecular compounds: made of 2 elements that share e- in a covalent bond
70
Explain why it is necessary to include a Roman numeral when naming an ionic compound with a transition metal.
You cannot determine the charge of a transition metal from the periodic table, so the Roman numeral is necessary to indicate the charge of the metal ion.
71
Give an example of a common compound and how the properties of the elements that make it up are different from that of the compound itself.
NaCl = sodium chloride
Sodium on its own is extremely explosive when in contact with water.
Chlorine on its own is extremely poisonous.
Together they make table salt, which we can eat.
72
Name the compound: Na2CO3
Sodium carbonate
73
Name the compound: NH4C2H3O2
ammonium acetate
74
Name the compound: BaCl2
barium chloride
75
Name the compound:Mn(OH)2
Manganese (II) hydroxide
76
Write the chemical formula for the following compounds: Magnesium fluoride
MgF2
77
Write the chemical formula for the following compounds: Calcium phosphate
Ca3(PO4)2
78
Write the chemical formula for the following compounds: Gold (II) carbonate
AuCO3
79
Write the chemical formula for the following compounds: Lead (IV) Sulfate
Pb(SO4)2
80
Explain what makes a metallic bond different from an ionic versus a covalent.
They are different from ionic because they do not lose their valence electrons, instead they share.
They are different from covalent because they do not result in a compound being formed.
81
Explain why the “sea of e-” forms in metallic bonds and why this makes them good electrical conductors.
The “sea of e-” forms because metals like to form lattices, but they are often bigger so their orbitals overlap. They also have lower electronegativities so their attraction for electrons in a bond is lower.
Because of this, the electrons are delocalized, making them good at creating a flow of e- when conducting electricity.
82
Describe the unique properties that result from metallic bonds.
Malleable (can be made into thin sheets)
Ductile (can be made into wires)
Can conduct electrical and thermal energy
Strong absorbers and reflectors of light
83
Give an example of an alloy. Include what metals make it up and how its structure aids in how the alloy is used.
14K gold is part gold, part silver and copper (usually). This alloy is stronger than gold on its own, making for more durable jewelry.
84
Explain the connection between electronegativity and the formation of a covalent bond.
When the electronegativity difference is small, and thus not big enough for one atom to take the electrons from the other, the atoms share instead in a covalent bond.
85
Differentiate between polar covalent and nonpolar covalent bonds. Give an example of each.
Polar covalent are when electrons are unequally shared.
Ex. In H2O, the oxygen atom has a stronger attraction and pull on the electrons than the hydrogen atoms, and thus the oxygen side is a little more negative and the hydrogen side a little more positive, creating poles.
Nonpolar covalent are when electrons are equally shared, and the difference in electronegativity is 0.
Ex. In O2, electrons are equally shared between the 2 atoms.
86
Differentiate between single, double, and triple bonds.
Single = 2 shared e-
Double = 4 shared e-
Triple = 6 shared e-
87
Explain how the VSEPR theory impacts how you draw Lewis structures for molecular compounds.
VSEPR theory is the tendency for electron pairs to be as far apart as possible from one another, and thus when you draw Lewis structures for molecular compounds, you should draw them in a way that spaces them out as much as you can.
88
Name the following compounds:
N2O3
Dinitrogen trioxide
89
Name the compound
SO2
sulfur dioxide
90
Name the compound:
N2S
dinoitrogen monosulfide
91
Name the compound:
B2H4
diboron tetrahydride
92
Write the chemical formula for the following compounds:
Boron tribromide
Bbr3
93
Write the chemical formula for the following compounds: Carbon tetrafluoride
CF4
94
Write the chemical formula for the following compounds: Diphosphorous trisulfide
P2S3
95
Write the chemical formula for the following compounds: Nitrogen trihydride
NH3
96
Differentiate between physical and chemical changes.
Physical changes do not impact a substance’s identity (Ex. state of matter change)
Chemical changes result in a new substance from a chemical reaction occurring (Ex. burning)
97
Explain the collision theory
Reacting particles must collide with enough energy for a chemical reaction to occur.
98
List several signs that can provide evidence that a chemical reaction has occurred.
Release of light or heat, sudden color change, odor change, gas released, formation of a precipitate
99
Explain the Law of Conservation of Mass and how it relates to chemical reactions.
Matter cannot be created or destroyed in a system – it can only change forms.
So even in a chemical reaction, the mass of the reactants should equal the mass of the products.
100
Differentiate between coefficients and subscripts in a chemical reaction.
Coefficients: are placed in front of a substance and represent ratios of reactants to products. They are used for balancing reactions and can be changed.
Subscripts: are the small numbers within each chemical formula that show the ratio of atoms in a compound. They cannot be changed because changing them changes the identity of the substance.
101
A+B=AB
Synthesis
102
2 of more reactants combine to form a new compound
Synthesis
103
When a substance reacts (burns) with O2
Combustion
104
Usually need heat or electricity to happen
decomposition
105
reacts with O2 oxide, metal and nonmetal oxide, salts
synthesis
106
1 reactant breaks down into 2 or more products
decomposition
107
AB = A + B
decomposition
108
reactant + O2 = product
combustion
109
when the reactant is a hydrocarbon, CO2 and H2O are products
combustion
110
1 element replaces a like element in a compound
single replacement
111
AB + CD = AD + CB
double replacement
112
the most reactive metals can react with H20 metal and hydroxide & H2
single replacement
113
often makes a precipitate, insoluble gas, or molecular compound
double replacement
114
A + BC = B + AC
single replacement
115
Classify the following reaction:
NF3 🡪 N2 + F2
decomposition
116
AlCl3 + K2SO4 🡪 KCl + Al2(SO4)3
double replacement
117
C4H10 + O2 🡪 CO2 + H2O
combustion
118
Pb(OH)2 + HCl 🡪 H2O + PbCl2
double replacement
119
4Al(s) + 3O2(g) 🡪 2Al2O3(s)
synthesis
120
Label each part of the neutralization reaction:
HCl(aq) + NaOH(aq) 🡪 NaCl(aq) + H2O(l)
compound 1 - acid
compound 2 - base
compound 3 - salt
compound 4 - water
121
Binary Ionic Compound:
made of 2 elements that transfer e- in an ionic bond
122
Compounds with Polyatomic Ions:
made of one (or more) covalently bonded charged group of atoms that transfer e- in an ionic bond.
123
Binary molecular compounds:
made of 2 elements that share e- in a covalent bond
