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Chemistry Honors > Chemistry Semester 2 Final Exam > Flashcards

Chemistry Semester 2 Final Exam Flashcards

1
Q

Specific patterns in the periodic table that make it possible to make predictions about an element’s properties

A

periodic trends

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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

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3
Q

One-half the distance between the nuclei of two identical atoms bonded together

A

atomic radius

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4
Q

A measure of how readily an atom will react

A

reactivity

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5
Q

When electrons in the inner energy levels block the attraction of the nucleus to the valence electrons

A

shielding effect

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6
Q

A measure of the ability of an atom in a bond to attract electrons

A

electronegativity

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7
Q

it is a measure of a neutral atom’s ability to gain an
electron

A

electron affinity

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8
Q

The energy required to remove one electron from a neutral atom of an element

A

ionization energy

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9
Q

the charge of an ion

A

oxidation numbers

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10
Q

The radius of a monatomic ion in an ionic compound

A

ionic radius

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11
Q

A force/mutual electrical attraction that holds atoms together in a substance

A

chemical bond

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12
Q

A measure of the ability of an atom in a bond to attract electrons

A

electronegativity

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13
Q

The tendency of atoms to prefer to have 8 valence electrons

A

Octet rule

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14
Q

Compound formed when 2 or more elements are covalently bonded

A

molecule

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15
Q

A binary ionic compound between 1 kind of metal atom and 1 kind of nonmetal atom

A

salt

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16
Q

Network of cations and anions that are mutually attracted to one another

A

crystal lattice

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17
Q

A compound made of positive and negative ions that chemically bond in a way that the charges equal out.

A

ionic compound

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18
Q

charged atoms

A

ion

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19
Q

Positively charged atom from losing electrons

A

cation

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20
Q

Negatively charged atom from gaining electrons

A

anion

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21
Q

A positively or negatively charged, covalently bonded group of atoms

A

polyatomic ion

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22
Q

Metals that “transition” because they can form many different ions

A

transition metal

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23
Q

An interaction that holds metal atoms together through a shared pool of valence electrons

A

metallic bond

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24
Q

Electrons not associated with a single atom or bond

A

delocalized electron

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25
Q

A solution of two metals mixed together

A

alloy

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26
Q

A compound formed when two or more atoms share electrons

A

covalent compound

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27
Q

The bond between nonmetallic atoms that share electrons

A

covalent bond

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28
Q

A neutral group of atoms held together by covalent bonds

A

molecule

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29
Q

When 2 atoms of the same element exist naturally as a molecule

A

diatomic elements

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30
Q

Process by which substances collide with enough energy that new bonds are formed between atoms, thus creating new substances

A

chemical reaction

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31
Q

starting substances

A

reactants

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32
Q

ending substances

A

products

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33
Q

dissolved in water

A

aqueous

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34
Q

A binary compound with at least 1 atom of oxygen

A

oxide

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35
Q

A compound made of cations and anions

A

salt

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36
Q

The decomposition of a substance by an electric current

A

electrolysis

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37
Q

A compound made of carbon and hydrogen

A

hydrocarbon

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38
Q

An insoluble solid that forms from the ions of two aqueous compounds

A

precipitate

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39
Q

A list of elements in the order in which they will easily undergo certain chemical reactions

A

activity series

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40
Q

Summarize the connection between Coulomb’s law and the periodic trends.

A

↑ 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.

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41
Q

Ionic radii:

A

Atoms that lose e- become smaller
Atoms that gain e- become larger
Tends to increase as you go DOWN a group

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42
Q

Explain how the shielding effect works.

A

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.

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43
Q

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.

A

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.

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44
Q

Identify the oxidation numbers for elements in groups 1, 2, and 13-17.

A

Group 1 🡪 +1
Group 2 🡪 +2
Group 13 🡪 +3
Group 14 🡪 +/- 4
Group 15 🡪 -3
Group 16 🡪 -2
Group 17 🡪 -1

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45
Q

Differentiate between the forces between different substances versus the forces within a substance.

A

Intermolecular forces are between different substances
Intramolecular forces are within the same substances (chemical bonds)

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46
Q

Explain what is meant by the statement, “Bonding is a spectrum.”

A

There aren’t just two types of bonds. It is a spectrum with ionic and nonpolar covalent at the two extremes.

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47
Q

Write a sentence to summarize the connections between chemical bonds, chemical reactions, and compounds.

A

Chemical bonds form through chemical reactions and result in new compounds.

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48
Q

Explain why most atoms form chemical bonds and why some atoms don’t.

A

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.

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49
Q

Explain the difference between what a chemical formula tells you about an ionic versus a covalent compound.

A

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.

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50
Q

IONIC OR COVALENT: electrons are transferred

A

ionic

51
Q

IONIC OR COVALENT: between metal and nonmetal

A

ionic

52
Q

IONIC OR COVALENT: always between ions

A

ionic

53
Q

IONIC OR COVALENT: can result in the formation of salt

A

ionic

54
Q

IONIC OR COVALENT: electrons are shared

A

covalent

55
Q

IONIC OR COVALENT: always between nonmetals

A

covalent

56
Q

IONIC OR COVALENT: can be polar or nonpolar

A

covalent

57
Q

IONIC OR COVALENT: results in the formation of a molecule

A

covalent

58
Q

IONIC OR COVALENT: crystalline solids

A

ionic

59
Q

IONIC OR COVALENT: high melting and boiling points

A

ionic

60
Q

IONIC OR COVALENT: can conduct electricity when dissolved in water

A

ionic

61
Q

IONIC OR COVALENT: can be solid, liquid, or gas

A

covalent

62
Q

IONIC OR COVALENT: low melting and boiling points

A

covalent

63
Q

IONIC OR COVALENT: cannot conduct electricity when dissolved in water

A

covalent

64
Q

List the number of atoms of each element in the compounds below. AlF3

A

Aluminum - 1
Fluorine - 3

65
Q

K3P

A

Potassium - 3
Phosphorous - 1

66
Q

Ca(ClO3)2

A

Calcium - 1
Chlorine - 2
Oxygen - 6

67
Q

Explain the connection between electronegativity and the formation of an ionic bond.

A

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
Q

Use an example of an ionic compound to explain the “rule of zero charge”.

A

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
Q

Differentiate between binary ionic compounds, compounds with polyatomic ions, and binary molecular compounds.

A

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
Q

Explain why it is necessary to include a Roman numeral when naming an ionic compound with a transition metal.

A

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
Q

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.

A

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
Q

Name the compound: Na2CO3

A

Sodium carbonate

73
Q

Name the compound: NH4C2H3O2

A

ammonium acetate

74
Q

Name the compound: BaCl2

A

barium chloride

75
Q

Name the compound:Mn(OH)2

A

Manganese (II) hydroxide

76
Q

Write the chemical formula for the following compounds: Magnesium fluoride

A

MgF2

77
Q

Write the chemical formula for the following compounds: Calcium phosphate

A

Ca3(PO4)2

78
Q

Write the chemical formula for the following compounds: Gold (II) carbonate

A

AuCO3

79
Q

Write the chemical formula for the following compounds: Lead (IV) Sulfate

A

Pb(SO4)2

80
Q

Explain what makes a metallic bond different from an ionic versus a covalent.

A

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
Q

Explain why the “sea of e-” forms in metallic bonds and why this makes them good electrical conductors.

A

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
Q

Describe the unique properties that result from metallic bonds.

A

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
Q

Give an example of an alloy. Include what metals make it up and how its structure aids in how the alloy is used.

A

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
Q

Explain the connection between electronegativity and the formation of a covalent bond.

A

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
Q

Differentiate between polar covalent and nonpolar covalent bonds. Give an example of each.

A

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
Q

Differentiate between single, double, and triple bonds.

A

Single = 2 shared e-
Double = 4 shared e-
Triple = 6 shared e-

87
Q

Explain how the VSEPR theory impacts how you draw Lewis structures for molecular compounds.

A

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
Q

Name the following compounds:
N2O3

A

Dinitrogen trioxide

89
Q

Name the compound
SO2

A

sulfur dioxide

90
Q

Name the compound:
N2S

A

dinoitrogen monosulfide

91
Q

Name the compound:
B2H4

A

diboron tetrahydride

92
Q

Write the chemical formula for the following compounds:
Boron tribromide

A

Bbr3

93
Q

Write the chemical formula for the following compounds: Carbon tetrafluoride

A

CF4

94
Q

Write the chemical formula for the following compounds: Diphosphorous trisulfide

A

P2S3

95
Q

Write the chemical formula for the following compounds: Nitrogen trihydride

A

NH3

96
Q

Differentiate between physical and chemical changes.

A

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
Q

Explain the collision theory

A

Reacting particles must collide with enough energy for a chemical reaction to occur.

98
Q

List several signs that can provide evidence that a chemical reaction has occurred.

A

Release of light or heat, sudden color change, odor change, gas released, formation of a precipitate

99
Q

Explain the Law of Conservation of Mass and how it relates to chemical reactions.

A

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
Q

Differentiate between coefficients and subscripts in a chemical reaction.

A

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
Q

A+B=AB

A

Synthesis

102
Q

2 of more reactants combine to form a new compound

A

Synthesis

103
Q

When a substance reacts (burns) with O2

A

Combustion

104
Q

Usually need heat or electricity to happen

A

decomposition

105
Q

reacts with O2 oxide, metal and nonmetal oxide, salts

A

synthesis

106
Q

1 reactant breaks down into 2 or more products

A

decomposition

107
Q

AB = A + B

A

decomposition

108
Q

reactant + O2 = product

A

combustion

109
Q

when the reactant is a hydrocarbon, CO2 and H2O are products

A

combustion

110
Q

1 element replaces a like element in a compound

A

single replacement

111
Q

AB + CD = AD + CB

A

double replacement

112
Q

the most reactive metals can react with H20 metal and hydroxide & H2

A

single replacement

113
Q

often makes a precipitate, insoluble gas, or molecular compound

A

double replacement

114
Q

A + BC = B + AC

A

single replacement

115
Q

Classify the following reaction:
NF3 🡪 N2 + F2

A

decomposition

116
Q

AlCl3 + K2SO4 🡪 KCl + Al2(SO4)3

A

double replacement

117
Q

C4H10 + O2 🡪 CO2 + H2O

A

combustion

118
Q

Pb(OH)2 + HCl 🡪 H2O + PbCl2

A

double replacement

119
Q

4Al(s) + 3O2(g) 🡪 2Al2O3(s)

A

synthesis

120
Q

Label each part of the neutralization reaction:
HCl(aq) + NaOH(aq) 🡪 NaCl(aq) + H2O(l)

A

compound 1 - acid
compound 2 - base
compound 3 - salt
compound 4 - water

121
Q

Binary Ionic Compound:

A

made of 2 elements that transfer e- in an ionic bond

122
Q

Compounds with Polyatomic Ions:

A

made of one (or more) covalently bonded charged group of atoms that transfer e- in an ionic bond.

123
Q

Binary molecular compounds:

A

made of 2 elements that share e- in a covalent bond

Chemistry Honors (18 decks)

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