Exam 2 Flashcards
(134 cards)
1
Q
Many main group metals _____ enough electrons to leave them with the same number of electrons as an atom of the preceding noble gas
A
lose
2
Q
Many non metals ____ enough electrons to give them the same number of electrons as an atom of the next noble gas
A
gain
3
Q
Ion Charge
Moving from far left to far right in the periodic table
A
positive charges of cations are equal to the group number
4
Q
Ion Charge
Moving from the far right to the far left in the periodic table
A
negative charges of anions are equal to the number of groups moved left from the noble gas, excluding the noble gas
5
Q
monoatomic ion
A
ions formed from only one atom
6
Q
polyatomic ions
A
electrically charged molecules (a group of bonded atoms with an overall charge)
7
Q
Oxyanions
A
polyatomic ions that contain one or more oxygen atoms
8
Q
Naming Oxyanions Rules
Nonmetal forms two oxyanions
A
-ate is the suffix used for the ion with the larger number of oxygen atoms
-ite is the suffix used for the ion with the smaller number of oxygen atoms
9
Q
Naming Oxyanions Rules
When a nonmetal forms more than two oxyanions
A
prefixes are used in addition to the -ate and -ite
per- (largest number of oxygens)
hypo- (smallest number of oxygens)
10
Q
Ionic bonds
A
electrostatic forces of attraction
when electrons are transferred, and ions form
11
Q
Covalent bond
A
when electrons are shared and molecules form
12
Q
Ionic compound
A
a compound that contains ions and is held together by ionic bonds
13
Q
metals and nonmetals form
A
ionic compounds
14
Q
Properties of ionic compounds
A
solids with high melting point and boiling points
nonconductive in solid form
conductive in molten form, poor conductor in solid form
dissolve readily in water
15
Q
The formula of an ionic compound must have
A
a ratio of ions such that the numbers of positive and negative charges are equal
16
Q
nonmetal and a nonmetal form a
A
covalent bond
17
Q
Molecular (covalent) compounds qualities
A
gases/liquids at room temp
insoluble in water
poor conductors of electricity
low- boiling
low-melting
18
Q
a cation is formed when
A
a neutral atom loses one or more electrons from its valence shell
19
Q
an anion forms when
A
a neutral atom gains one or more electrons in its valence shell
20
Q
Bond length
A
determined by distance which lowest potential energy is achieved
21
Q
Potential energy
A
if atoms continue to approach each other, the positive charges of the nuclei start repelling each other, increasing potential energy
22
Q
What is needed to break chemical bonds
A
energy
23
Q
breaking chemical bonds is considered
A
endothermic
24
Q
forming chemical bonds
A
releases energy
25
Forming chemical bonds is considered
exothermic
26
pure covalent bond
when atoms forming a covalent bond are identical, the electrons in the bond are shared equally
27
when the atoms linked by a covalent bond are different, the bonding electrons are
shared but not equally
shifts electron density towards the atom that is more attractive to electrons
28
electronegativity
measures the tendency of an atom to attract electrons towards itself
29
nonpolar versus polar is determined by
electronegativity
30
the more strongly an atom attracts electrons the larger its
electronegativity
31
electrons in a polar covalent bond are
shifted toward the more electronegative atom
32
Trend of electronegativity
increases from left to right across a period and decreases down a group
33
most electronegative element
fluorine
34
Electron affinity
a measurable quantity of the energy released or absorbed when an isolated gas-phase atom acquired an electron
35
How do we measure polarity expected in a bond
the absolute value of the difference in electronegativity of two bonded atoms
nonpolar- small
polar- large
36
As electronegativity difference increases between two atoms the bond becomes
more ionic
37
Nomenclature
a collection of rules for naming things
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Ionic compound naming rules
1. name the cation first, followed by the name of the anion
(monoatomic cation is just given the name of the element)
2. a monoatomic anion is given the name of the element with its ending replaced by the suffix -ide
4. a polyatomic ion is just given the name of the ion
Ionic compounds examples
NaCl- sodium chloride
KBr- potassium bromide
Polyatomic ionic compound examples
KC2H3O2- potassium acetate
NH4Cl- ammonium chloride
39
Naming Ionic compounds containing a metal with a variable charge
Charge of the metal ion is specified by a Roman numeral in parentheses after the name of the metal
examples
FeCl2- iron(II) chloride
FeCl3- iron(III) chloride
40
Naming ionic hydrates
1. name the anhydrous compound (per usual rules)
2. add the word hydrate with a Greek prefix denoting the number of water molecules
examples
CuSO4 5H2O- Copper (II) Sulfate pentahydrate
41
Hydrate
compound, often ionic, that contains one or more water molecules bound within its crystals
42
Prefix for 1
mono
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Prefix for 2
di
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prefix for 3
tri
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prefix for 4
tetra
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prefix for 5
penta
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prefix for 6
hexa
48
prefix for 7
hepta
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prefix for 8
octa
50
prefix for 9
nona
51
prefix for 10
deca
52
Rules for naming Molecular compounds
the name of the more metallic element is named first
the name of the nonmetallic is next with the suffix =ide
the numbers of each element are designated by Greek prefixes
examples
SO2- sulfur dioxide
N2O4- dinitrogen tetroxide
53
Naming binary acids
1. the word hydrogen is changed to the prefix hydro-
2. the other metallic element name is modified by adding the suffix- ic
3. the word acid is added as a second word
example
HF, hydrogen fluoride- hydrofluoric acid
HCl, hydrogen Chloride- hydrochloric acid
54
Oxyacids
compounds that contain hydrogen, oxygen, and at least one other element, are bonded in such a way as to impart acidic properties to the compound
55
Naming oxyacids
1. omit hydrogen
2. start with the root name of the anion
3. replace -ate with -ic and -ite with -ous
4. add acid
examples
HC2H3O2 anion name- acetate= acetic acid
HClO4 anion name perchlorate= perchloric acid
56
Lewis symbol
consists of an elemental symbol surrounded by one dot for each of its valence electrons
used to describe valence electron configurations of atoms and monatomic ions
formation of covalent bonds are shown as well
57
When forming a cation or anion in a lewis structure
the number of electrons does not change
58
lone pairs
electrons that are not used in bonding
59
single bond
a single shared pair of electrons
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octet rule
the tendency of main group atoms to form enough bonds to obtain eight valence electrons
61
hydrogen needs how many electrons to fill its shell
two
62
exceptions to octet rule
hydrogen- needs two
transition metals
inner transition elements
63
double bond
two pairs of electrons are shared between atoms
64
triple bond
three pairs of electrons are shared by a pair of atoms
65
Drawing lewis structures procedures
1. determine the total number of valence electrons
- for cations subtract one electron for each positive charge
-for anions add one electron for each negative charge
2. draw a skeleton structure, arranging the atoms around a central atom
- least electronegative element should be central
-connect each atom to the central atom with a single bond
3. distribute the remaining electrons on the terminal atoms (except hydrogen) completing the octet rule for each atom
4. place all remaining electrons on the central atom
5. rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible
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Odd-electron molecules
have an odd number of valence electrons, and therefore have an unpaired electron
67
Electron-deficient molecules
have a central atom that has fewer than needed for a noble gas configuration
68
hypervalent molecules
have a central atom that has more electrons than needed for a noble gas configuration (more than 4)
69
Free radicals
molecules that contain an odd number of electrons
70
Elements in the second period can accommodate how many electrons
8 because they have 4 valence orbitals
71
Elements in the third and higher periods have
more than 4 valence orbitals and can share more than 4 pairs of electrons with other atoms
72
Formal charge
the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms
73
Formal charge equation
valence shell electrons- # of lone pair electrons- 1/2 # bonding electrons
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molecular structure
the arrangement of atoms in a molecule or ion
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Steps for molecular structure
1. molecular structure in which all formal charges are zero is preferred
2. if the lewis structure must have nonzero formal charges, the arrangement with the smallest nonzero formal charge is preferred
3. lewis structures are preferable when adjacent formal charges are zero or of the opposite sign
4. when we must choose among several lewis structures with similar distributions of charge, the structure with the negative formal charges on the more electronegative atom is preferable
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Resonance
if two or more lewis structures with the same arrangement of atoms can be written for a molecule or ion
77
Resonance forms
each individual lewis structure of a resonance
78
Resonance hybrid
the actual electronic structure of the molecule
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bond angle
the angle between any two bonds that include a common atom, usually measured in degrees
80
bond distance
the distance between the nuclei of two bonded atoms along the straight line joining the nuclei
81
Valence Shell Electron-pair Repulsion Theory
enables us to predict the molecular structure, including approximate bond angles around a central atom, of a molecule from an examination of the number of bonds and lone electron pairs in its lewis structure
82
Region of electron density
a single bond, a double bond, a triple bond, or a lone pair each count as one region of electron density
83
Electron-pair geometry
geometry that includes all electron pairs
84
Molecular structure indicates
only the placement of the atoms in the molecule
85
Order of repulsions an order of the amount of space occupied by different kinds of electron pairs
lone pair-lone pair> lone pair- bonding pair>bonding pair- bonding pair
86
Sizes from largest to smallest electron densities
lone pair>triple bond>double bond>single bond
87
Axial position
if we hold the model of a trigonal bipyramid by the two axial positions, we have an axis around which we can rotate the model
88
Equatorial position
three positions form an equator around the middle of the molecule
89
Predicting electron pair geometry and molecular structure steps
1. write the lewis structure
2. count the number of regions of electron density around the central atom
3. identify the electron pair geometry based on the number of regions of electron density
4. use the number of lone pairs to determine the molecular structure
90
Bond dipole moment
the measure of polarity of a chemical bond between two atoms in a molecule
91
electrons are pulled toward the (more or less) electronegative atom
more
92
dipole moment formula
dipole moment=separated charge times distance between
93
To be polar a moleculemust
contain at least one polar covalent bond
have a molecular structure such that the sum of the vectors of each bond dipole moment does not cancel
94
Valence bond theory
electrons reside in orbitals
covalent bonds form when electrons are shared by the overlapping of half filled atomic orbitals (each containing a single electron)
Electrons in these overlapping orbitals must have opposite spins
to form a covalent bond an atom must have an unpaired electron
number of bonds formed is determined by the number of unpaired electrons
95
Sigma bonds are produced by
1. two s orbitals
2. an s orbital and a p orbital
3. two p orbitals
96
Sigma bond
a covalent bond in which the electron density is concentrated in the region along the internuclear axis
97
Pi Bond
a type of covalent bond that results from the side-by-side overlap of two Pi orbitals
98
node
the plane with no probability of finding and electron
99
All single bonds are
sigma bonds
100
double bond consists of
one sigma bond and one pi bond
101
triple bond consists of
one sigma bond and two pi bonds
102
Two atomic orbitals can hybridize
one s orbital+ one p orbital=
two sp hybrid orbitals
103
three atomic orbitals can hybridize
one s orbital+ two p orbitals=
three sp2 hybrid orbitals
occurs when a central atom is surrounded by three regions of electron density
104
four atomic orbitals can hybridize
one s orbital+3 p orbitals=
four sp3 hybrid orbitals
occurs when a central atom is surrounded by four regions of electron density
105
Five atomic orbitals can hybridize
one s orbital+ 3 p orbitals+ one d orbital
five sp3d hybrid orbitals
occurs when a central atom is surrounded by five regions of electron density
106
Six atomic orbitals can hybridize
one s orbital+3 p orbitals+ two d orbitals=
6 sp3d2 orbitals
occurs when a central atom is surrounded by six regions of electron density
107
Assignment of hybrid orbitals on central atoms
1. determine lewis structure
2. determine the number of regions of electron density around an atom using VSEPR theory
3. assign the set of hybridized orbitals that corresponds to its geometry
108
paramagnetism
when we pour liquid oxygen past a strong magnet it collects between the poles of the magnet and defies gravity
arises in molecules that have unpaired electrons
109
paramagnetic samples will
appear heavier
the number of unpaired electrons is calculated based on increased weight
110
Molecular bond theory
describes the distribution of electrons in molecules so that distribution of electrons in atoms is described using atomic orbitals
111
molecular orbital
region of space in which a valence electron in a molecule is likely to be found
112
Linear combination of atomic orbitals
the mathematical process of combining atomic orbitals is to generate these orbitals
113
bonding orbitals
adding electrons creates a force that holds the 2 nuclei together
114
antibonding orbitals
the attractive force between the nuclei and these electrons pulls the two nuclei apart
115
Molecular orbital diagram
shows relative energy levels of atomic and molecular orbitals
116
bond order
number of bonding electrons- number of nonbinding electrons all divided by 2
117
Formula mass
sum of the average atomic masses of all the atoms in the substances formula
118
molecular mass
formula mass- another name for molecule not ions
119
percent composition
the percentage by mass of each element in a compound
120
determination of empirical formula
1. convert element masses to moles using molar masses
2. divide each number of moles by the smallest number of moles
3. if necessary, multiply by an integer, to give the smallest whole number ratio of subscripts
121
determination of empirical formula from its percent composition
1. convert percent composition to masses of elements by assuming a 100g sample of the compound
2. convert element to moles using molar mass
3. divide each number of moles by the smallest number of moles
4. if necessary, multiply by an integer, to give the smallest shole-number ratio of subscripts
122
Molecular formula
molecular or molar mass divided by empirical formula
123
solutions or homogenous mixtures
uniform composition and properties throughout its entire volume
124
concentration
the relative amount of a given solution component
125
Solvent
component with a concentration that is significantly greater than that of all other components
126
solute
component that is present at a much lower concentration
127
aqueous solution
a solution in which water is the solvent
128
Molarity
the number of moles of solute is exactly 1 L of the solution
129
dilution
the process whereby the concentration of a solution is lessened by the addition of solvent
130
Stock solution
more concentrated
131
Mass percentage
is defined as the ratio of the component's mass to the solution's mass, expressed as a percentage
Mass percentage= mass of component divided by mass of solution all times 100
132
Volume percentage
a ratio of a solute's mass to the solution's volume expressed as a percentage
133
parts per million
mass solute/mass of solution X 10^6
134
Parts per billion
Mass solute/mass of solution X 10^9
