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Chemistry 1320 > Exam 2 > Flashcards

Exam 2 Flashcards

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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

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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

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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

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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

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5
Q

monoatomic ion

A

ions formed from only one atom

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6
Q

polyatomic ions

A

electrically charged molecules (a group of bonded atoms with an overall charge)

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7
Q

Oxyanions

A

polyatomic ions that contain one or more oxygen atoms

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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

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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)

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10
Q

Ionic bonds

A

electrostatic forces of attraction
when electrons are transferred, and ions form

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11
Q

Covalent bond

A

when electrons are shared and molecules form

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12
Q

Ionic compound

A

a compound that contains ions and is held together by ionic bonds

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13
Q

metals and nonmetals form

A

ionic compounds

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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

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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

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16
Q

nonmetal and a nonmetal form a

A

covalent bond

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17
Q

Molecular (covalent) compounds qualities

A

gases/liquids at room temp
insoluble in water
poor conductors of electricity
low- boiling
low-melting

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18
Q

a cation is formed when

A

a neutral atom loses one or more electrons from its valence shell

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19
Q

an anion forms when

A

a neutral atom gains one or more electrons in its valence shell

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20
Q

Bond length

A

determined by distance which lowest potential energy is achieved

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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

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22
Q

What is needed to break chemical bonds

A

energy

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23
Q

breaking chemical bonds is considered

A

endothermic

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24
Q

forming chemical bonds

A

releases energy

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25
Q

Forming chemical bonds is considered

A

exothermic

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26
Q

pure covalent bond

A

when atoms forming a covalent bond are identical, the electrons in the bond are shared equally

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27
Q

when the atoms linked by a covalent bond are different, the bonding electrons are

A

shared but not equally
shifts electron density towards the atom that is more attractive to electrons

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28
Q

electronegativity

A

measures the tendency of an atom to attract electrons towards itself

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29
Q

nonpolar versus polar is determined by

A

electronegativity

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30
Q

the more strongly an atom attracts electrons the larger its

A

electronegativity

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31
Q

electrons in a polar covalent bond are

A

shifted toward the more electronegative atom

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32
Q

Trend of electronegativity

A

increases from left to right across a period and decreases down a group

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33
Q

most electronegative element

A

fluorine

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34
Q

Electron affinity

A

a measurable quantity of the energy released or absorbed when an isolated gas-phase atom acquired an electron

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35
Q

How do we measure polarity expected in a bond

A

the absolute value of the difference in electronegativity of two bonded atoms
nonpolar- small
polar- large

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36
Q

As electronegativity difference increases between two atoms the bond becomes

A

more ionic

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37
Q

Nomenclature

A

a collection of rules for naming things

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38
Q

Ionic compound naming rules

A
  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
  3. 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

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39
Q

Naming Ionic compounds containing a metal with a variable charge

A

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

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40
Q

Naming ionic hydrates

A
  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
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41
Q

Hydrate

A

compound, often ionic, that contains one or more water molecules bound within its crystals

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42
Q

Prefix for 1

A

mono

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43
Q

Prefix for 2

A

di

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44
Q

prefix for 3

A

tri

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45
Q

prefix for 4

A

tetra

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46
Q

prefix for 5

A

penta

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47
Q

prefix for 6

A

hexa

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48
Q

prefix for 7

A

hepta

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49
Q

prefix for 8

A

octa

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50
Q

prefix for 9

A

nona

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51
Q

prefix for 10

A

deca

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52
Q

Rules for naming Molecular compounds

A

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

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53
Q

Naming binary acids

A
  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
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54
Q

Oxyacids

A

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
Q

Naming oxyacids

A
  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
Q

Lewis symbol

A

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
Q

When forming a cation or anion in a lewis structure

A

the number of electrons does not change

58
Q

lone pairs

A

electrons that are not used in bonding

59
Q

single bond

A

a single shared pair of electrons

60
Q

octet rule

A

the tendency of main group atoms to form enough bonds to obtain eight valence electrons

61
Q

hydrogen needs how many electrons to fill its shell

A

two

62
Q

exceptions to octet rule

A

hydrogen- needs two
transition metals
inner transition elements

63
Q

double bond

A

two pairs of electrons are shared between atoms

64
Q

triple bond

A

three pairs of electrons are shared by a pair of atoms

65
Q

Drawing lewis structures procedures

A
  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
66
Q

Odd-electron molecules

A

have an odd number of valence electrons, and therefore have an unpaired electron

67
Q

Electron-deficient molecules

A

have a central atom that has fewer than needed for a noble gas configuration

68
Q

hypervalent molecules

A

have a central atom that has more electrons than needed for a noble gas configuration (more than 4)

69
Q

Free radicals

A

molecules that contain an odd number of electrons

70
Q

Elements in the second period can accommodate how many electrons

A

8 because they have 4 valence orbitals

71
Q

Elements in the third and higher periods have

A

more than 4 valence orbitals and can share more than 4 pairs of electrons with other atoms

72
Q

Formal charge

A

the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms

73
Q

Formal charge equation

A

valence shell electrons- # of lone pair electrons- 1/2 # bonding electrons

74
Q

molecular structure

A

the arrangement of atoms in a molecule or ion

75
Q

Steps for molecular structure

A
  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
76
Q

Resonance

A

if two or more lewis structures with the same arrangement of atoms can be written for a molecule or ion

77
Q

Resonance forms

A

each individual lewis structure of a resonance

78
Q

Resonance hybrid

A

the actual electronic structure of the molecule

79
Q

bond angle

A

the angle between any two bonds that include a common atom, usually measured in degrees

80
Q

bond distance

A

the distance between the nuclei of two bonded atoms along the straight line joining the nuclei

81
Q

Valence Shell Electron-pair Repulsion Theory

A

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
Q

Region of electron density

A

a single bond, a double bond, a triple bond, or a lone pair each count as one region of electron density

83
Q

Electron-pair geometry

A

geometry that includes all electron pairs

84
Q

Molecular structure indicates

A

only the placement of the atoms in the molecule

85
Q

Order of repulsions an order of the amount of space occupied by different kinds of electron pairs

A

lone pair-lone pair> lone pair- bonding pair>bonding pair- bonding pair

86
Q

Sizes from largest to smallest electron densities

A

lone pair>triple bond>double bond>single bond

87
Q

Axial position

A

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
Q

Equatorial position

A

three positions form an equator around the middle of the molecule

89
Q

Predicting electron pair geometry and molecular structure steps

A
  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
Q

Bond dipole moment

A

the measure of polarity of a chemical bond between two atoms in a molecule

91
Q

electrons are pulled toward the (more or less) electronegative atom

A

more

92
Q

dipole moment formula

A

dipole moment=separated charge times distance between

93
Q

To be polar a moleculemust

A

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
Q

Valence bond theory

A

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
Q

Sigma bonds are produced by

A
  1. two s orbitals
  2. an s orbital and a p orbital
  3. two p orbitals
96
Q

Sigma bond

A

a covalent bond in which the electron density is concentrated in the region along the internuclear axis

97
Q

Pi Bond

A

a type of covalent bond that results from the side-by-side overlap of two Pi orbitals

98
Q

node

A

the plane with no probability of finding and electron

99
Q

All single bonds are

A

sigma bonds

100
Q

double bond consists of

A

one sigma bond and one pi bond

101
Q

triple bond consists of

A

one sigma bond and two pi bonds

102
Q

Two atomic orbitals can hybridize
one s orbital+ one p orbital=

A

two sp hybrid orbitals

103
Q

three atomic orbitals can hybridize
one s orbital+ two p orbitals=

A

three sp2 hybrid orbitals
occurs when a central atom is surrounded by three regions of electron density

104
Q

four atomic orbitals can hybridize
one s orbital+3 p orbitals=

A

four sp3 hybrid orbitals
occurs when a central atom is surrounded by four regions of electron density

105
Q

Five atomic orbitals can hybridize
one s orbital+ 3 p orbitals+ one d orbital

A

five sp3d hybrid orbitals
occurs when a central atom is surrounded by five regions of electron density

106
Q

Six atomic orbitals can hybridize
one s orbital+3 p orbitals+ two d orbitals=

A

6 sp3d2 orbitals
occurs when a central atom is surrounded by six regions of electron density

107
Q

Assignment of hybrid orbitals on central atoms

A
  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
Q

paramagnetism

A

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
Q

paramagnetic samples will

A

appear heavier
the number of unpaired electrons is calculated based on increased weight

110
Q

Molecular bond theory

A

describes the distribution of electrons in molecules so that distribution of electrons in atoms is described using atomic orbitals

111
Q

molecular orbital

A

region of space in which a valence electron in a molecule is likely to be found

112
Q

Linear combination of atomic orbitals

A

the mathematical process of combining atomic orbitals is to generate these orbitals

113
Q

bonding orbitals

A

adding electrons creates a force that holds the 2 nuclei together

114
Q

antibonding orbitals

A

the attractive force between the nuclei and these electrons pulls the two nuclei apart

115
Q

Molecular orbital diagram

A

shows relative energy levels of atomic and molecular orbitals

116
Q

bond order

A

number of bonding electrons- number of nonbinding electrons all divided by 2

117
Q

Formula mass

A

sum of the average atomic masses of all the atoms in the substances formula

118
Q

molecular mass

A

formula mass- another name for molecule not ions

119
Q

percent composition

A

the percentage by mass of each element in a compound

120
Q

determination of empirical formula

A
  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
Q

determination of empirical formula from its percent composition

A
  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
Q

Molecular formula

A

molecular or molar mass divided by empirical formula

123
Q

solutions or homogenous mixtures

A

uniform composition and properties throughout its entire volume

124
Q

concentration

A

the relative amount of a given solution component

125
Q

Solvent

A

component with a concentration that is significantly greater than that of all other components

126
Q

solute

A

component that is present at a much lower concentration

127
Q

aqueous solution

A

a solution in which water is the solvent

128
Q

Molarity

A

the number of moles of solute is exactly 1 L of the solution

129
Q

dilution

A

the process whereby the concentration of a solution is lessened by the addition of solvent

130
Q

Stock solution

A

more concentrated

131
Q

Mass percentage

A

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
Q

Volume percentage

A

a ratio of a solute’s mass to the solution’s volume expressed as a percentage

133
Q

parts per million

A

mass solute/mass of solution X 10^6

134
Q

Parts per billion

A

Mass solute/mass of solution X 10^9

Chemistry 1320 (2 decks)

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